c 


MEDICAL 


COLLEGE  OF  PHARMACY 


maoy 


0&2!fcrn!a  Co!'        ;f  Pharmacy 


AN  INTRODUCTION 

TO 

CHEMICAL  ANALYSIS 


ROCKWOOD 


AN  INTRODUCTION 


TO 


CHEMICAL  ANALYSIS 

FOR  STUDENTS  OF  MEDICINE,  PHARMACY 
AND  DENTISTRY 


BY 

ELBERT  W.  j^QCKWOOD,  M.  D.,  PH.  D. 

PROFESSOR  OF  CHEMISTRY  AND  TOXICOLOGY  AND  HEAD  OF  THE  DEPARTMENT 
CHEMISTRY  IN  THE  UNIVERSITY  OF  IOWA,  AUTHOR  OF  "A  LABORATORY 
MANUAL  OF  PHYSIOLOGICAL  CHEMISTRY  " 


California  College  of  Pharmacy 


FOURTH  REVISED  EDITION 
WITH  20  ILLUSTRATIONS 


PHILADELPHIA 

P.  BLAKISTON'S  SON  &  CO. 

1012  WALNUT  STREET 


COPYRIGHT,  1913,  BY  P.  BLAKISTON'S  SON  &  Co. 


y  Q  K  K 


D75 


PREFACE  TO  FOURTH    EDITION 


ADVANTAGE  has  been  taken  of  the  exhaustion  of  the  last 
edition  to  make  such  alterations  as  appeared  desirable  without 
materially  changing  the  original  plan  and  scope  of  the  book- 
that  is,  to  leave  it  an  introductory  guide  for  the  first-year 
college  student.  The  changes  are  chiefly  the  simplification  of 
methods  where  experience  has  shown  that  the  student  is 
most  likely  to  have  difficulty,  together  with  the  correction  of 
a  fe,w  errors.  For  suggestions  I  have  been  indebted  to  in- 
structors in  this  department,  particularly  to  Mr.  J.  E.  Booge. 

THE  UNIVERSITY  OF  IOWA. 


PREFACE  TO  THIRD  EDITION 


WITHIN  the  last  few  years  the  methods  and  theories  of 
physical  chemistry  have  been  more  and  more  successfully 
applied  in  different  fields  of  medicine.  Among  the  theories 
which  have  been  most  fruitful  in  results  stands  prominently  the 
ionic  theory  of  the  nature  of  solutions  and  the  action  of  many 
classes  of  dissolved  compounds.  It  is  self-evident  that  a 
knowledge  of  this  is  a  necessity  for  those  who  wish  to  acquire 
an  acquaintance  with  chemistry  from  the  present-day  stand- 
point or  to  keep  pace  with  its  developments.  The  surest 
means  of  gaining  familiarity  with  these  conceptions  is  by 
applying  them  practically  in  every-day  work. 

In  the  present  revision  of  this  book,  therefore,  greater 
prominence  than  before  has  been  given  to  the  ionic  explana- 
tions of  analytical  reactions.  This  has  been  done  partly 
because  of  the  clearness  with  which  such  explanations  can  be 
made  and  partly  to  enable  the  student  to  assimilate  the 
general  theory.  At  the  same  time  it  has  not  been  thought 
best  to  wholly  abandon  the  older  terminology  which  is  still 
largely  used  in  medical  science. 

The  arrangement  of  the  work  has  otherwise  been  changed 
but  very  little.  No  attempt  has  been  made  to  produce  a 
complete  manual  of  qualitative  analysis,  but  rather  to  pro- 
vide such  an  introduction  to  the  subject  as  may  show  its 
application  to  other  branches  of  science  and  perhaps  lead  the 
student  to  a  desire  for  a  more  extended  course. 

THE  UNIVERSITY  OF  IOWA. 


vii 


PREFACE  TO  FIRST  EDITION 


ALTHOUGH  it  may  be  desirable  that  elementary  chemistry 
should  be  completed  before  the  professional  college  is  entered, 
this  is  impractical  in  case  of  many  American  schools.  The 
value  of  analytical  chemistry  in  stimulating  observation, 
power  of  discrimination,  independence  and  self-reliance  has 
long  been  recognized,  as  well  as  its  services  in  affording  an 
easy  introduction  to  chemical  work.  It  may,  however,  be 
carried  on  as  a  handicraft,  without  being  of  assistance  in 
demonstrating  the  fundamentals  of  chemistry.  The  medical 
student  not  infrequently  regards  it  as  of  no  value  except  for 
the  purpose  of  making  analyses  and  for  these  he  believes  he 
will  have  little  use.  Though  the  importance  of  chemical 
analysis  per  se  will  probably  always  be  slight,  its  points  of 
contact  with  the  other  subjects  in  the  medical  curriculum  are 
so  numerous  that,  when  rightly  carried  on,  it  not  only  stimu- 
lates interest  in  its  own  pursuit,  but  gives  a  mastery  over  other 
branches  which  can  be  as  easily  attained  in  no  other  way. 

The  chemical  incompatibility  of  medicinal  substances  is  so 
intimately  related  to  their  chemical  reactions  that  for  their 
intelligent  use  a  thorough  acquaintance  with  the  latter  is  a 
prerequisite.  The  same  is  true  of  the  chemical  antidotes  for 
poisons.  For  this  reason  the  reactions  are  quite  fully  given 
here.  The  properties  and  methods  of  manufacture  of  many 
substances  employed  for  medical  or  dental  purposes  are  also 
illustrated  by  the  analytical  reactions,  which  can  therefore 
be  used  to  impress  them  upon  the  mind  of  the  student  and  thus 
make  a  connected  chain  of  what  is  often  learned  as  dismem- 
bered facts. 

ix 


X  PREFACE    TO    FIRST    EDITION 

It  is  natural  that  chemical  analysis  for  beginners  should 
be  differently  conducted  with  professional  students  than  with 
those  who  do  not  desire  to  apply  it  to  any  particular  branch  of 
knowledge.  This  book  has  been  arranged  for  students  of 
medicine,  pharmacy,  and  dentistry,  much  of  whose  territory 
is  common.  It  assumes  that  some  study  has  been  devoted  to 
general  chemistry  or  that  this  is  a  contemporaneous  course. 
It  is  designed  to  furnish  a  scientific  basis  for  more  technical 
courses  but  not  to  supplant  these,  and  to  give  the  familiarity 
with  chemicals  and  manipulative  methods,  which  is  so  neces- 
sary for  real  success  in  some  lines  of  medical  work. 

There  is  no  intention  to  make  of  the  student  an  analytical 
chemist  or  mere  mechanical  manipulator.  In  many  cases  the 
work  is  abridged,  as  in  the  detection  of  poisons,  where  only 
the  principal  ones  are  considered  or  those  which  best  illustrate 
the  methods  of  such  analysis.  At  the  same  time  no  attempt 
has  been  made  so  far  to  cut  down  the  work  that,  with  a  mere 
smattering  of  knowledge,  the  student  finishes,  believing  him- 
self competent  to  meet  all  problems  that  he  may  encounter. 

To  accomplish  the  purpose  outlined,  series  of  questions 
have  been  inserted.  The  answers  to  these  may  be  found 
partly  in  the  experimental  work  previously  done,  and  partly 
through  reading  in  other  departments  of  chemistry  or  of 
medicine.  Others  will  suggest  themselves  to  the  instructor, 
and  it  is  only  by  insisting  upon  such  outside  study  that  the 
greatest  value  can  be  gained  from  such  a  course  as  this.  If 
followed  out  they  are  a  stimulus  to  individual  application,  they 
prevent  mechanical  working  without  thinking,  and,  by  con- 
necting chemical  analysis  with  general  chemistry ,  materia  med  - 
ica,  physiology,  toxicology,  andotherdepartmentsof  medicine, 
they  help  to  make  clear  the  unity  of  the  complete  course. 

But  few  equations  have  been  given,  and  those  usually 
only  the  more  difficult  ones.  These  benefit  the  student 
only  when  he  can  write  them  for  himself  and  he  should  do 


PREFACE    TO    FIRST   EDITION  XI 

this  as  far  as  possible  in  the  time  allowed.  For  a  similar 
reason  no  tables  are  given  for  finding  without  labor  the  results 
of  volumetric  analyses.  To  represent  the  metric  system  as 
something  more  than  a  theory  and  to  prepare  for  its  future 
practical  use  all  measurements  are  stated  in  metric  denomi- 
nations. Degrees  of  temperature  are  given  by  the  centigrade 
thermometer. 

While  the  length  of  the  medical  course  often  forbids  any 
extended  work  in  quantitative  analysis,  some  practical  work 
in  this  is  indispensable,  partly  in  order  to  familiarize  the  stu- 
dent with  principles,  partly  because  of  its  applications  to  other 
departments  of  medicine.  Volumetric  methods  are  admirably 
adapted  for  both  these  purposes  and  enough  are  given  to  il- 
lustrate the  more  common  and  to  indicate  how  they  may  be 
extended.  They  include  the  preparation  of  the  standard 
solutions,  as  well  as  the  use  of  these,  to  that  the  student 
may,  if  necessary,  be  in  condition  to  undertake  t  he  whole  proc- 
ess in  the  practice  of  his  profession  and  not  be  reduced  to  a 
state  of  helplessness  if  the  emergency  should  arise.  For  the 
same  reason,  in  a  special  table  is  included  the  preparation  and 
testing  of  the  qualitative  reagents. 

To  make  the  course  more  interesting,  by  showing  some  of 
its  applications,  chapters  are  added  on  the  testing  of  water, 
the  detection  of  poisons,  and  analysis  by  means  of  the  blow- 
pipe. While  the  latter  is  of  subordinate  importance  for 
students  of  medicine  and  pharmacy,  it  has,  in  this  laboratory, 
proved  itself  of  value  in  demonstrating  to  students  of  dentis- 
try the  physical  and  chemical  properties  of  the  metals  and 
their  alloys  in  a  manner  not  possible  by  wet  methods  of 
testing.  The  length  of  the  course  can  be  modified  to  fit  the 
curriculum  by  omitting  the  less  important  parts  or  by  varying 
the  number  of  unknown  substances  to  be  analyzed  under  the 
separate  divisions  of  the  subject. 

THE  UNIVERSITY  OF  IOWA. 


TO  THE  STUDENT 

1.  PERFORM  no  operation  without  a  reason.     Ask  yourself 
in  advance  the  object,  and  afterward  what  has  occurred; 
for  example,  in  precipitating  or  washing — what  is  removed? 
what  remains? 

2.  Do  your  own  work;  use  your  own  judgment,  and  let 
your  neighbor  do  the  same. 

3.  Avoid  the  use  of  an  excess  of  materials  or  reagents.     Add 
the  reagent  slowly,  as  much  as  is  necessary  and  no  more.     Do 
not  use  concentrated  acids  unless  they  are  specific  ally  calledfor . 

4.  When  a  strongly  acid  solution,  or  one  which  gives  an 
offensive  odor  is  to  be  boiled,  or  when  any  acid  solution  is  to 
be  evaporated,  do  this  under  a  hood  that  the  gases  may  not 
remain  in  the  room. 

5.  Except  for  cleaning,  always  use  distilled  water. 

6.  Always  use  pure  chemicals  for  reagents,  but  never  be 
sure  that  they  are  so  without  proving  them. 

7.  Never  put  platinum  wire,  stirring  rod  or  other  object 
into  a  reagent  bottle.     Do  not  return  to  the  bottle  any 
reagent  that  has  been  removed. 

8.  In  using  the  reagents  never  lay  the  stopper  down. 
Hold  it  between  the  second  and  third  fingers  and  replace  it 
immediately  in  the  bottle. 

9.  Do  not  throw  into  the  sink  concentrated  acids,  strong 
solutions  of  mercury,  or  solid  refuse  like  broken  glass  or  filter- 
paper.     Put  them  in  the  waste  jar. 

10.  Too  great  care  cannot  be  exercised  as  to  cleanliness. 
Ha,ve  a  cloth  or  towel  and  keep  apparatus  and  desk  in  good 
order.    Each  student  will  be  held  responsible  for  the  condi- 
tion in  which  his  desk  is  left. 

xiii 


TABLE  OF  CONTENTS 


INTRODUCTION i 

PART  I 

QUALITATIVE  ANALYSIS 

CHAPTER     I.— Metals  (Cations) 23 

CHAPTER    II. — Acids  (Anions) 96 

CHAPTER  III — Organic  Compounds 122 

PART  II 

VOLUMETRIC  ANALYSIS 

CHAPTER  *  I. — General  Principles      148 

CHAPTER    II. — Analysis  by  Neutralization 158 

CHAPTER  III. — Analysis  by  Oxidation  and  Reduction 167 

CHAPTER  IV. — Analysis  by  Precipitation] 177 

PART  III 

APPLIED  ANALYSIS 

CHAPTER     I. — The  Sanitary  Examination  of  Water 185 

CHAPTER    II. — The  Detection  of  Poisons 203 

CHAPTER  III. — Analysis  by  Means  of  the  Blowpipe 221 

.,    PART  IV 

The  Preparation  and  Testing  of  Reagents 231 

The  Chemical  Elements— Symbols  and  Atomic  Weights 240 

The  Metric  System 241 

INDEX 243 


xv 


CHEMICAL  ANALYSIS 


INTRODUCTION 

THE  object  of  chemical  analysis  is  the  determination  of  the 
chemical  composition  of  matter.  This  determination  may  be 
either  of  the  kind  of  its  components  or  of  their  amounts. 
The  methods  used  for  the  former  belong  to  qualitative  anal- 
ysis; for  the  latter  to  quantitative. 

In  qualitative  analysis  the  substance  studied  is  not  only 
examined  alone,  but  it  is  subjected  to  the  action  of  certain 
forces  or  of  chemical  compounds  called  reagents.  These 
are  designed  to  make  more  evident  the  physical  properties  of 
the  substance  or  to  produce  characteristic  compounds  or 
changes.  All  such  phenomena,  called  the  reactions  of  the 
substance,  can  be  used  in  its  identification.  Before  entering 
upon  the  practical  work  of  qualitative  chemical  analysis 
some  explanation  is  demanded  relating  to  the  apparatus 
and  reagents  most  in  use,  to  the  common  processes,  and  to 
some  of  the  terms  employed. 

Solutions 

In  the  majority  of  cases,  although  not  always,  the  qualita- 
tive tests  are  made  upon  substances  in  solution.  A  solution 
is  formed  wfren  a  solid,  liquid,  or  gas  is  taken  up  by  a  liquid 
so  that  it  loses  its  usual  physical  properties  and  can  no 
longer  be  perceived.  Solutions  are  regarded  as  of  two  kinds, 
physical  and  chemical. 

Physical  solutions  are  similar  to  those  which  result  from 
mixing  salt  or  sugar  with  water.  There  is  no  apparent 


INTRODUCTION    TO    CHEMICAL   ANALYSIS 

chemical  change  in  the  substance  dissolved,  which  is  left 
unaltered  when  the  solvent  evaporates.  In  physical  solu- 
tions of  solids  in  liquids  a  rise  in  temperature,  as  a  rule,  in- 
creases not  only  the  rapidity  of  solution,  but  also  the  amount 
of  the  solid  capable  of  being  dissolved.  With  solutions  of 
gases  in  liquids  the  opposite  is  true  and  most  dissolved  gases 
will  be  driven  out  of  solution  by  an  increased  temperature. 

Chemical  solutions  are  such  as  the  one  formed  when  metals 
are  acted  upon  by  an  acid.  The  metals  disappear  but  unite 
with  the  solvent  so  that  when  the  latter  evaporates  an  entirely 
different  substance  from  the  original  one  remains.  In  this 
case  heating  lessens  the  time  of  solution,  but  does  not  neces- 
sarily increase  the  amount  of  metal  dissolved.  In  chemical 
solutions  of  gases,  however,  the  gas  is  often  expelled  by  heat- 
ing, the  compound  being  thus  decomposed. 

A  concentrated  solution  is  one  in  which  the  solvent  contains 
a  large  amount  of  the  dissolved  substance  (the  solute) ;  a  di- 
lute solution,  one  in  which  the  amount  dissolved  is  small. 
When  the  solvent  contains  as  much  of  the  solute  as  it  can  take 
up  it  is  said  to  be  saturated.  A  solution  which  has  been  satu- 
rated with  a  solid  at  a  high  temperature  ordinarily  deposits  a 
part  of  the  solid  if  the  temperature  falls.  A  few,  as  the  tem- 
perature is  decreased,  will  hold  what  has  gone  into  solution, 
until  the  liquid  is  shaken  or  some  solid  matter  is  introduced, 
whereupon  a  large  quantity  of  the  dissolved  material  sepa- 
rates from  the  liquid.  Solutions  containing  this  excessive 
amount  are  said  to  be  super-saturated.  Where  a  solid 
remains  floating  in  a  liquid  without  dissolving  it  is  said  to  be 
suspended  or  in  suspension. 

Most  solids  which  have  been  dissolved  in  a  liquid  and 
which  afterward  separate  through  a  change  in  temperature 
or  a  decrease  in  the  volume  of  the  solvent  are  crystalline  in 
form.  Such  a  method  of  preparation  is  called  crystallization. 
A  crystal  is  a  solid  which  has  a  regular  form  bounded  by  plane 


INTRODUCTION  3 

surfaces,  the  angles  between  these  being  constant  for  the 
same  species  of  crystal.  Large  crystals  are  produced  only 
when  they  form  slowly.  It  has  been  found  that  impure  solids 
after  having  been  dissolved  in  water  or  other  fluid  and  then 
allowed  to  crystallize  leave  most  of  their  impurities  in  the 
liquid  or  "  mother-liquor  "  so  that  this  affords  one  of  the  best 
methods  for  purifying  such  substances. 

The  Heating  of  Liquids 

For  making  solutions  and  for  heating  liquids  the  chemist 
commonly  uses  test-tubes,  beakers  or  flasks  of  glass,  or  dishes 
of  porcelain.  Test-tubes  are  made  of  thin  glass  and  are 


FIG.  i. — Apparatus  for  solution,  i.  Test-tubes  in  wooden  support.  2. 
Beakers.  3.  Iron  stand  with  adjustable  rings  for  supporting  objects  while 
they  are  heated.  On  one  of  the  rings  is  a  sand-bath.  4.  Flasks. 


designed  to  be  heated  in  contact  with  the  flame  without  break- 
ing. The  flame  should,  however,  not  be  allowed  to  touch  the 
tube  above  the  liquid,  since  it  easily  becomes  superheated 
there  and  breaks  when  moistened  later.  The  heat  should 


4  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

not  be  applied  to  the  bottom  of  the  tube  alone,  as  the  steam 
bubbles  thus  produced  might  suddenly  force  the  hot  liquid 
from  the  tube,  but  all  the  parts  of  the  liquid  should  be  heated 
by  moving  and  turning  the  tube  in  the  flame.  On  account  of 
this  same  danger  the  examinations  of  the  boiling  substance 
should  be  made  from  the  side  and  never  from  above,  and  a 
test-tube  while  boiling  should  never  be  pointed  toward 
another  person.  A  tube  half -full  can  be  brought  to  the  boil- 
ing point  while  it  is  held  in  the  fingers,  although  holders  are 
sometimes  used.  If  one  is  desired  it  can  be  made  from  a 
strip  of  paper,  which  permits  the  rotation  of  the  tube. 

If  long  boiling  is  requisite  a  beaker  is  to  be  preferred  to  a 
test-tube.  As  these  are  somewhat  thicker  on  the  sides  than 
at  the  bottom,  they  are  not  so  well  adapted  to  being  heated  in 
contact  with  the  flame.  They  may  be  supported  by  an  iron- 
wire  gauze  or  plate,  which  can  be  heated  to  a  low  redness 
without  danger  to  the  beaker  if  it  contains  a  liquid.  As  with 
any  other  glass  vessels  they  should  not  be  heated  above  the 
liquid.  Instead  of  the  gauze  or  plate,  a  sheet  of  asbestos  is 
sometimes  used  or  a  sand-bath — that  is,  a  shallow  iron  dish 
with  a  thin  layer  of  sand  which  distributes  the  heat  evenly  to 
the  vessel.  The  two  latter,  although  they  are  safe,  waste 
more  of  the  heat  than  does  the  gauze.  Glass  flasks  can  be 
used  in  the  same  way  as  the  beakers.  Evaporation  is  less 
from  these  than  from  beakers,  which  is  often  desirable.  All 
glass  vessels  composed  of  the  ordinary  German  glass  are  some- 
what attacked  by  boiling  water,  and  especially  so  by  alkaline 
solutions.  If  great  accuracy  is  desired  those  made  of  a  re- 
sistant glass  should  be  substituted  for  the  former.  Dishes  of 
porcelain  are  but  little  affected  by  the  ordinary  reagents  even 
at  the  boiling-point,  and  the  danger  of  their  breaking  is  much 
less.  A  further  advantage  is  that  they  may  be  heated  with 
the  naked  flame.  Liquids  heated  in  them  for  a  long  time, 
however,  suffer  a  considerable  loss  by  evaporation. 


INTRODUCTION  5 

Evaporation 

Evaporation  is  the  process  by  which  a  volatile  substance 
is  separated  from  a  less  volatile  solid  or  liquid,  the  more 
volatile  compound  being  allowed  to  escape.  It  may  take  place 
slowly  at  the  ordinary  temperature,  as,  for  example,  when 
tinctures  become  more  concentrated,  through  loss  of  alcohol*, 
by  standing  in  open  bottles.  It  is  hastened  by  the  higher 


FIG.  2. — Apparatus  for  evaporation,  i.  An  extemporized  steam-bath — a 
beaker  of  water  on  which  a  dish  can  be  heated.  2.  A  copper  steam-bath  with 
tube  supplying  water  and  allowing  the  excess  to  escape,  thus  maintaining  a 
constant  level.  The  funnel  above  excludes  dust.  3.  A  sulphuric  acid  vacuum- 
desiccator  connected  with  air  pump.  In  front  are  evaporating  dishes  of 
porcelain,  platinum  and  glass. 

temperatures  and  is  most  frequently  carried  on  by  this  means 
in  chemical  operations.  Thus  it  is  made  use  of  in  removing 
from  solutions  an  excessive  amount  of  a  volatile  acid  like 
hydrochloric,  or  even  of  one  which  is  less  readily  converted  to  a 
gas,  like  sulphuric,  in  which  case  to  effect  its  removal  the  liquid 


6  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

must  be  evaporated  to  dry  ness.  When  a  liquid  is  to  be  evap- 
orated a  wide  and  shallow  dish  should  be  used  as  the  vapors 
most  easily  pass  off  from  this.  From  a  test-tube,  on  the 
other  hand,  evaporation  is  slow,  the  vapors  condensing  above 
to  the  liquid  form,  and  flowing  back  into  the  tube.  It  is  often 
advisable  to  conduct  the  process  at  some  definite  temperature, 
for  instance  not  above  ioo01  in  order  to  avoid  burning  or  de- 


FIG.  3. — Apparatus  for  distillation.  A  flask  connected  with  a  Liebig's  con- 
denser and  provided  with  a  thermometer  for  showing  the  boiling-point  of  the 
liquid  which  it  contains. 

composition.  This  may  be  done  by  placing  the  liquid  to  be 
evaporated  in  a  dish  and  setting  this  on  a  vessel  in  which 
water  is  boiling,  known  as  a  water-bath  or  steam-bath.  For 
small  operations  a  beaker  will  answer  for  the  lower  vessel. 

1  In  this  work  all  temperatures  are  given  in  the  centigrade  scale,  where  the 
freezing-point  of  water  is  at  o°  and  the  boiling-point  100°. 


INTRODUCTION 


All  evaporations  of  injurious  or  offensive  gases  should  be  con- 
ducted in  a  hood  or  fume  chamber — an  enclosure  connected 
with  a  ventilating  shaft.  Aqueous  solutions  can  be  evapo- 
rated at  low  temperatures  by  placing  the  dish  over  a  vessel  of 
some  substance  which  has  an  affinity  for 
moisture,  like  concentrated  sulphuric 
acid,  and  placing  over  the  whole  an  air- 
tight cover.  This  is  called  a  desiccator. 
The  process  is  much  hastened  by  ex- 
hausting the  air  from  the  apparatus. 

Distillation 

Distillation  is  similar  to  evaporation  ex- 
cept that  the  expelled  volatile  substance 
does  not  escape  but  is  collected  by  being 
again  condensed  to  a  liquid  through  cool- 
ing. This  is  commonly  effected  by  pass- 
ing the  vapors  into  a  tube  which  is  sur- 
rounded by  circulating  cold  water.  If 
a  thermometer  is  suspended  with  its 
bulb  in  the  vapor  during  distillation  it 
indicates  the  boiling-point  of  the  liquid. 
This,  with  the  melting-point,  is  of  the 
greatest  value  in  the  identification  of 
many  compounds,  especially  the  organic 
ones.  The  melting-point  is  determined 
after  heating  a  short,  small  glass  tube  to 

redness  in  the  middle,  then  drawing  it  of  the  melting-point  of 

.      ?         a  solid, 
out  to  a  very  small  diameter.     A  piece 

of  this  closed  at  the  lower  end  holds  some  of  the  powder  be- 
side a  thermometer-bulb  in  a  liquid,  like  water  or  sulphuric 
acid.  By  heating  gradually  until  the  powder  melts,  then 
reading  the  temperature,  its  melting-point  is  ascertained. 


8  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Precipitation 

Precipitation  is  the  process  of,  changing  dissolved  sub- 
stances into  the  insoluble  state.  It  may  be  without  chemical 
change,  as  when  the  dissolved  matters  of  tinctures  are  pre- 
cipitated by  water,  simply  because  they  are  insoluble  in  the 
latter.  Or,  on  the  other  hand,  the  precipitate  may  be  a  new 
chemical  compound  which  has  been  produced  by  the  mutual 
decomposition  of  two  dissolved  substances,  as  when  red  mer- 
curic iodid  is  formed  by  mixing  a  solution  of  potassium  iodid 
with  one  of  corrosive  sublimate.  B  oth  varieties  of  precipitates 
are  important  in  medicine  and  pharmacy,  but  the  latter 
will  be  the  ones  most  frequently  illustrated  in  qualitative 
analysis.  They  are  usually  heavier  than  the  surrounding  liquid 
and  consequently  settle  to  the  bottom,  but  some  of  the  floccu- 
lent  kinds  float  for  a  long  time  before  sinking.  Precipitation 
is  valuable  as  a  means  of  identifying  an  unknown  metal,  since 
by  this  means  characteristic  compounds  often  appear;  it  is 
used  to  separate  one  substance  from  another  by  converting 
the  one  into  an  insoluble  form;  it  is  also  of  use  in  the  prepara- 
tion and  purification  of  medicinal  as  well  as  other  substances. 
Thus  by  the  addition  of  a  soluble  carbonate  to  a  calcium  solu- 
tion " precipitated  chalk,"  a  very  pure  form  of  calcium  car- 
bonate, is  produced.  When  used  in  separation  the  filtrate 
must  always  be  tested  with  more  of  the  reagent  in  order  to  be 
certain  that  precipitation  is  complete. 

Filtration 

A  precipitate  or  other  undissolved  solid  can  be  separated 
from  a  liquid  in  which  it  is  suspended  by  filtration.  To  ac- 
complish this  the  mixture  is  poured  upon  some  porous  mate- 
rial which  allows  the  fluid  to  pass  through  but  retains  the 
solids.  A  porous  paper  is  the  most  common  filtering  agent. 
It  is  supported  in  a  glass  funnel  of  such  an  angle  that  if  a  cir- 


INTRODUCTION  9 

cular  piece  of  the  paper  is  folded  in  the  center  and  then  again 
in  the  center  at  right  angles  to  the  first  fold  it  will  when 
opened  make  a  cone-shaped  paper  funnel  which  fits  the  glass 
exactly.  To  obtain  the  most  rapid  and  clean  filtration  it 
should  be  made  to  adhere  to  the  sides  of  the  funnel  by  mois- 
tening it  with  fluid  of  the  same  kind  as  that  to  be  filtered  and 
pressing  it  against  the  glass.  It  ought  to  be  of  such  a  size  as 
not  to  reach  quite  to  the  rim  of  the  funnel.  The  only  ones  of 
the  common  reagents  which  attack  the  cellulose  of  which  the 
paper  is  composed  are  concentrated  acids  or  strong  solutions 
of  sodium  hydroxid  or  potassium  hydroxid.  The  cheaper 
filter-papers  contain  compounds  of  the  metals,  such  as  cal- 
cium and  iron,  and  these  may  be  dissolved  by  reagents 
which  are  being  filtered,  thereby  rendering  the  latter  impure. . 
Where  great  accuracy  is  important  the  paper  is,  therefore, 
previous  to  being  used,  washed  with  hydrochloric  or  hydro- 
fluoric acid  to  remove  impurities. 

When  it  is  desired  to  obtain  the  solution — called  the  fil- 
trate— as  quickly  as  possible  without  saving  the  precipitate, 
the  plaited  filter  can  be  employed.  This  is  made  by  folding 
double  as  before,  and  then  into  eight  or  sixteen  folds,  bending 
the  paper  each  time  in  the  same  direction.  Then  each  divi- 
sion is  folded  in  the  opposite  way,  without  pressing  at  the  tip, 
to  avoid  breaking,  so  that  the  filter  looks  like  a  closed  fan. 
When  it  is  opened  the  hollow  cone  has  fluted  sides,  giving 
twice  the  surface  of  the  plain  filter.  A  plug  of  absorbent 
cotton  loosely  placed  in  the  funnel  is  also  convenient  for 
obtaining  the  filtrate  when  the  precipitate  is  to  be  discarded. 
Cotton  consists,  like  the  paper,  of  cellulose.  Liquids  which 
decompose  cellulose  should  be  filtered  through  asbestos  or 
glass  wool. 

Many  precipitates  are  so  gelatinous  that  they  clog  the 
filter  and  the  filtration  proceeds  very  slowly.  By  passing  the 
stem  of  a  funnel  through  a  rubber  stopper  into  a  vessel  in 


10 


INTRODUCTION   TO    CHEMICAL   ANALYSIS 


which  a  partial  vacuum  can  be  created,  the  solution  can  be 
more  rapidly  drawn  through  the  filter.  The  air  may  be  ex- 
hausted from  the  vessel  below  by  means  of  a  filter  pump  or 
aspirator  attached  to  the  water  faucet  of  the  laboratory.  To 
avoid  breaking  the  paper  its  tip  must  then  be  supported  by 
inserting  below  it  in  the  funnel  a  cone  of  platinum,  muslin  or 
some  other  strong  material  or  a  specially  toughened  paper 
must  be  employed.  It  should  be  remembered  that  a  hot 
solution  generally  filters  more  rapidly  than  a  cold  one. 


FIG.  5. — A  pparatusfor  filtration,  i .  A  funnel  fitted  to  filtering  a  flask  which 
is  connected  with  a  pump  for  the  production  of  a  vacuum  and  consequently 
an  increase  in  the  rapidity  of  the  filtration.  2.  A  support  holding  funnels  for 
filtration.  3.  Washing  bottles.  In  front  are  packages  of  filter-paper  with  two 
plaited  filters. 

After  the  precipitate  has  been  collected  on  the  filter  it  is 
usually  necessary  to  further  purify  it  by  removing  the  liquid 
with  which  it  is  saturated  and  which  contains  soluble  matter. 
This  is  done  by  washing,  generally  using  for  this  purpose 
distilled  water,  the  hot  being  ordinarily  preferable  to  the  cold. 
The  water  may  be  poured  on  from  any  vessel,  but  is  most 


INTRODUCTION  1 1 

conveniently  applied  from  a  washing-bottle.  This  is  con- 
structed of  a  bottle  with  a  stopper  perforated  with  two  holes. 
Through  one  a  long,  bent  tube  passes  to  the  bottom.  It  is 
contracted  to  a  narrow  opening  at  the  outer  end  and  serves 
for  the  exit  of  a  jet  which  is  forced  out  by  blowing  into  the 
other  opening  through  a  short  tube.  By  such  a  small  stream 
the  precipitate  can  be  thoroughly  mixed  with  the  water  and 
if  so  desired,  can  be  rinsed  from  the  paper.  For  cold  water 
any  thick  glass  bottle  of  convenient  size  will  serve,  but  if  it  is 
to  be  heated  a  thin  bottomed  flask  must  be  employed. 
Washing  with  hot  water  is  more  effectual  than  with  cold.  To 
ascertain  whether  the  washing  is  complete  a  drop  of  the 
filtrate  can  be  evaporated  on  a  platinum  foil  and  the  amount 
of  dissolved  matter  observed. 

Sometimes  when,  in  spite  of  the  above  methods,  the  filtra- 
tion is  very  slow,  it  may  be  better  to  wash  by  decantation — 
that  is,  by  letting  the  precipitate  settle,  pouring  off  the  liquid 
without  disturbing  the  solid  and  repeating  as  many  times  as 
is  necessary.  With  heavy  compounds  like  those  of  mercury 
this  can  be  done  very  rapidly. 

Chemical  Changes  at  High  Temperatures 

In  order  to  produce  certain  chemical  changes,  or  in  the 
preparation  of  some  substances,  a  higher  degree  of  heat  is  re- 
quired than  that  of  boiling  water.  Even  fusion  or  melting 
may  be  necessary.  In  many  qualitative  tests  where  small 
quantities  of  material  are  used  the  substance  may  be  sup- 
ported on  a  piece  of  charcoal,  porcelain  or  platinum.  The 
first  necessitates  the  aid  of  the  blowpipe;  the  second  breaks 
too  easily;  the  third  is  preferable  because,  while  it  is  itself 
affected  by  few  reagents  and  there  is  no  danger  of  breaking, 
it  permits  a  high  heat  to  be  attained.  A  pair  of  forceps  or 
tongs  will  support  it  long  enough  in  the  flame  to  accomplish 


12 


INTRODUCTION   TO   CHEMICAL   ANALYSIS 


the  fusion.  With  large  quantities  of  material,  or  where  the 
heating  must  be  long-continued  a  crucible  can  be  advantage- 
ously substituted.  These  are  most  commonly  made  of 
platinum,  porcelain,  or  clay. 

Platinum  is  attacked  by  chlorin  or  by  any  mixtures,  like 
aqua  regia,  which  produce  it.  This  is  true  not  only  during 
fusion,  but  from  solutions  at  low  temperatures.  It  will  also 
be  affected  by  heating  with  fusible  metals,  since  these  form 
alloys  with  it.  Reducible  compounds  of  such  metals  mixed 


FIG.  6. — Apparatus  for  fusion,  i.  Porcelain  crucible.  2.  Hessian  or  clay 
crucibles  with  a  pipe-stem  triangle  leaning  against  the  middle  one.  On  this 
the  crucible  can  be  supported  while  it  is  heated.  3.  A  graphite  crucible.  4. 
A  platinum  crucible  with  cover.  In  front  are  forceps  and  crucible  tongs. 

with  carbon  should  not  be  heated  in  platinum  nor  should  the 
fusible  sulphids  or  the  caustic  alkalies,  all  of  which  combine 
with  it. 

In  the  absence  of  illuminating  gas,  the  alcohol  lamp  will 
serve  as  a  source  of  heat  since  it  does  not  smoke  cold  objects 
held  in  the  flame,  while  giving  a  great  deal  of  heat.  With 
gas,  which  is  more  convenient  as  well  as  cheaper  than  alcohol, 
the  Bunsen  burner  is  used.  This  consists  essentially  of  a 
base,  from  which  the  gas  is  delivered  by  a  small  opening,  and 


INTRODUCTION  13 

above  this  a  tube  with  openings  near  the  base  through  which 
is  admitted  the  air  which  then  mixes  with  the  gas.  This 
mixing  brings  about  the  complete  combustion  of  the  carbon 
compounds  of  the  gas,  none  being  left  to  be  deposited  as  soot, 
a  high  temperature  being  thereby  produced.  The  openings 
below  should  be  so  regulated  as  to  allow  entrance  to  enough 
air  to  change  the  color  of  the  flame  from  a  yellow  to  a  light 
blue.  More  than  this  lowers  the  temperature.  With  an  ex- 


FIG.  7. — Sources  of  heat.  i.  Two  forms  of  Bunsen  burners.  2.  A  cluster  of 
five  Bunsen  burners  for  greater  heat.  3.  A  Bunsen  burner  with  iron  chimney 
to  shut  off  drafts  of  air.  4.  A  blast  lamp  in  position  to  heat  the  platinum 
crucible  above.  5.  A  gas  furnace  for  more  intense  heat  with  larger  crucibles. 
In  front  are  two  forms  of  blowpipes. 

cess  of  air,  too,  the  flame  often  " strikes  back"  or  burns  within 
the  tube  at  the  bottom,  changing  in  color  first  to  a  yellow,  after- 
ward to  green.  In  this  case  the  gas  must  be  turned  off,  the 
size  of  the  air  vents  reduced  and  the  burner  relighted. 

For  higher  degrees  of  heat  the  blast-lamp  will  be  advanta- 
geous. This  is  similar  to  the  Bunsen  burner,  except  that  the 
air  is  forced  in  through  one  tube  while  the  gas  enters  by 
another,  the  two  mixing  just  before  they  are  burned.  A 
bellows  or  water-blast  will  furnish  enough  air  for  this.  For 


14  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

large  crucibles  there  is  less  waste  of  heat  if  the  crucible  is  sur- 
rounded by  a  fire-clay  box  or  furnace  and  the  flame  from  the 
blast-lamp  is  forced  in  below. 

For  the  better  examination  of  the  products  solid  substances 
are  occasionally  heated  or  fused  in  a  glass  tube  closed  at  one 
end,  thus  resembling  a  small  test-tube.  It  is  made  from  a 
piece  of  straight  glass  tubing  about  12  to  15  cm.  in  length  and 
with  a  bore  of  about  4  mm.  If  this  is 
heated  in  the  middle  until  it  is  thoroughly 
softened,  turning  continually,  it  can  be 
drawn  apart  while  in  the  flame  into  two 
such  tubes.  The  long  point  can  be  re- 
moved by  touching  it  while  red  hot  with 
another  piece  of  glass.  It  should  be  used 
for  only  one  fusion. 

The  flame  of  the  Bunsen  burner  is  not 
only  a  valuable  agent  as  a  heat-producer, 
but  also  for  the  production  of  chemical 
changes.  If  the  holes  below  are  closed  so 
that  no  air  enters,  the  flame  is  luminous 
but  smoky,  depositing  a  coating  of  un- 
burned  carbon  or  soot  upon  a  cold  object 
held  in  it.  The  temperature  is  then  not 
very  high  because  of  the  incomplete  com- 
FIG  8  —The  flame  bustion.  As  air  is  admitted  the  yellow 
of  the  Bunsen  burner  color  disappears  and  the  flame  becomes 

showing  the  four  zones.    .  . 

blue.  There  is  an  increase  in  tempera- 
ture and  objects  heated  are  not  smoked.  An  examination 
shows  this  flame  to  be  composed  of  several  parts,  (a)  At  the 
base,  within,  is  a  blue  cone  varying  in  height  with  the  amount 
of  gas.  This  is  a  mixture  of  air  and  gas  which  is  not  being 
burned.  If  a  splinter  is  thrust  into  it  and  held  there  it  will 
burn  off  at  the  margin  of  the  flame  before  the  ignition  of  the 
inner  end  occurs.  (6)  This  cone  is  outlined  by  a  line  of  lighter 


INTRODUCTION  15 

blue.  Here  combustion  is  proceeding  but  not  complete,  there 
being  present  unburned  carbon  compounds  of  the  gas .  Under 
these  conditions  they  will  reduce  oxygen  compounds,  that  is, 
remove  the  oxygen  from  them  (reduction  or  deoxidation). 
This  part  of  the  flame  is  called  the  reducing  or  deoxidizing 
zone.  A  copper  wire  oxidized  by  holding  it  above  the  flame 
until  it  is  black  will  lose  its  oxygen  and  its  dark  color  after  it  is 
held  a  few  seconds  in  it.  (c)  Outside  the  zone  of  reduction  lies 
a  wider  zone,  the  hottest  part  of  the  flame,  since  in  it  combus- 
tion is  complete.  It  is  called  the  zone  of  fusion  and  should  be 
used  when  heat  alone  is  desired,  (d)  The  outer  margin  of  the 
flame  is  bluish- white  and  rather  indistinct.  There  the  air  is 
in  excess  and  it  is  consequently  able  to  oxidize,  or  give  oxygen 
to,  metals  or  other  substances  heated  in  it.  It  is  called, 
therefore,  the  zone  of  oxidation  or  oxiding  zone. 

In  a  similar  manner  the  mouth  blowpipe  is  made  use  of  in 
effecting  chemical  changes.  It  is  merely  a  bent  tube  of  metal 
with  a  small  opening  through  which  air  can  be  blown  into  a 
flame  and  the  action  of  the 
latter  can  be  thereby  modi- 
fied. The  flame  to  be  used 
is  the  yellow  one  made  by 
closing  the  holes  of  the  Bun- 
sen  burner  or  by  slipping  into 

it     a     smaller     tube    with    a   FIG.  9.— Tne  oxidizing  flame  produced 

narrow  opening  above.    This 

latter  is  preferable  since  less  gas  will  be  necessary  and  the 
flame  can  be  better  directed.  To  obtain  an  oxidizing  action 
the  yellow  flame  should  be  about  5  cm.  high  (2  inches) 
and,  with  the  tip  of  the  blowpipe  just  within  the  flame, 
enough  air  should  be  used  to  make  it  blue,  directing  the 
blast  through  the  greatest  diameter  of  the  flame.  With 
the  blowpipe  also  the  blue  flame  should  be  employed  when 
great  heat  is  desired.  For  the  reducing  flame  perhaps 


1 6  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

50  per  cent,  more  gas  should  be  used;  the  blowpipe  should 
be  held  outside  the  flame  and  only  enough  air  be  forced^in  to 
make  the  flame  horizontal,  not  sufficient  to  change  the  yellow 


FIG.  10. — The  reducing  flame  produced  by  the  blowpipe  (color,  yellow). 

color  to  a  blue.  The  reducing  action  is  due  here  as  in  the 
Bunsen-burner  flame  to  the  incompletely  oxidized  hot 
carbon. 

Chemical  Reagents 

Since  chemical  action  proceeds  more  readily  between  dis- 
solved compounds  than  solids,  the  reagents  are  in  most  cases 
used  in  solution.  Unless  for  some  particular  reason  large 
amounts  of  these,  and  concentrated  solutions,  should  be 
avoided.  The  dilute  acids,  as  a  rule,  give  better  results  than 
the  concentrated.  When  the  latter  are  required  such  direc- 
tions will  be  given.  In  separating  two  or  more  metals  an 
" excess"  of  the  reagent  must  be  used,  but  this  does  not  imply 
an  excessively  large  volume  or  weight.  It  means  merely 
enough  to  do  the  work  for  which  it  is  added.  If  this  is  precipi- 
tation enough  must  be  used  to  completely  remove  the  precipi- 
table  elements  and  the  solution  must  afterward  be  tested  with 
more  of  the  reagent  to  make  sure  that  this  has  been  done. 
If  it  is  to  acidify  or  make  alkaline  a  solution  the  reagent 
should  be  slowly  poured  in  and,  after  mixing,  tests  should  be 
made  to  determine  whether  the  liquid  has  the  proper  reaction. 

The  reaction  of  a  solution  or  other  substance  is  ascertained 


INTRODUCTION  17 

by  its  effect  upon  colored  compounds,  the  one  most  commonly 
in  use  being  litmus-paper — that  is  paper  colored  blue,  red  or 
purple  with  a  vegetable  coloring  matter — litmus.  If  the 
liquid  turns  this. red  it  is  said  to  have  an  acid  reaction;  if  it 
turns  blue  the  reaction  is  alkaline.  It  is  possible  to  make  use 
of  other  colored  substances  besides  this  and  a  number  will  be 
30  used  in  this  course  of  study.  The  bases  which  are  soluble 
in  water  change  the  color  of  red  litmus  to  blue,  or  are  alka- 
line in  reaction;  the  insoluble  bases  do  not  affect  the  color. 
Soluble  acids  redden  litmus-paper,  but  an  acid  reaction  in  a 
liquid  does  not  necessarily  prove  that  a  free  acid  is  present, 
for  other  compounds  can  produce  the  same  result.  An  acid  is 
a  compound  of  hydrogen  with  one  or  more  negative  (non- 
metallic)  elements;  and  in  the  acids  the  hydrogen  can  be  dis- 
placed by  metals  or  similar  substances  and  thus  form  a  salt. 

Zn+H2SO4  =  H2+ZnSO4 

Salts  are  therefore  seen  to  be  composed  of  a  metal  or  a 
similar  substance  united  with  the  negative  part  of  an  acid. 
They  can  also  be  formed  by  the  action  of  an  acid  upon  a  base, 
which  latter  consists  of  a  metal,  or  a  group  of  elements  of 
similar  properties,  combined  with  OH.  In  such  cases  water 
is  produced  at  the  same  time. 

NaOH-f  HCl  =  NaCl-r-H20,  or 
2NH4OH+H2S04=  (NH4)2S04+2H20 

The  Theory  of  Solution  (Ionic  Theory) 

A  knowledge  of  many  chemical  changes,  including  a  major- 
ity of  those  of  qualitative  analysis,  is  facilitated  by  a  knowledge 
of  the  ionic  theory,  a  brief  outline  of  which  is  given  here.  Ac- 
cording to  this  theory,  in  aqueous  solutions  such  are  commonly 
used  in  analytical  chemistry,  acids,  bases,  and  salts  undergo  a 
partial  decomposition  into  their  positive  and  negative  com- 


1 8  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

ponents,  which  are  then  called  ions.  This  change  is  known  as 
dissociation  or  ionization.  The  degree  of  ionization  increases 
with  the  dilution,  but  varies  for  different  compounds,  being 
greater  for  strong  acids  and  bases,  or  salts  of  these,  than  for 
the  weak  ones.i  The  ions  differ  from  the  chemical  elements  in 
that  they  do  not  separate  from  the  solution  in  such  a  way  as  to 
be  visible.  They  may  be  composed  of  a  single  form  of  matter 
such  as  the  ions  of  H  or  Cl  from  hydrochloric  acid,  or  they 
may  contain  two  or  more  forms  of  matter,  such  as  NH4  of  am- 
monium salts  or  C2H3O2  of  acetic  acid.  All  electrolytes,  as 
compounds  are  called  whose  solutions  conduct  the  electric 
current,  are  ionizable.  The  ions  of  a  solution  are  set  in  motion 
if  an  electric  current  is  passed  through  it  and  are  named  from 
the  electrodes  toward  which  they  travel,  cations  going  in  the 
direction  of  the  cathode  or  negative,  and  anions  toward  the 
anode  or  positive  electrode. 

Under  these  conditions  the  ions  can  be  considered  as  carriers 
of  electricity.  The  univalent  ones  have  a  single  charge  and 
those  of  higher  valence  have  electric  charges  proportional  to 

the  valence,  the  nature  and  amount  being  indicated  by  signs, 

+       ++     +++  - 

such  as  Na,  Ba,  Al,  Cl,  NOs,  P04.  Instead  of  the  plus  and 
minus  signs  these  charges  are  often  indicated  by  dots  and 
dashes',  as  Na',  SCU",  etc.  When  these  signs  are  present  it 
means  that  the  substances  are  not  in  the  form  of  elements,  but 
of  ions.  Thus  the  properties  of  H  and  H',  of  Cl  and  Cl',  of  S 
and  S",  are  entirely  different.  When  the  ion  loses  its  electric 
charge  it  becomes  an  element  or  group  of  elements. 

The  properties  of  dilute  aqueous  solutions  are  largely  those 
of  their  ions;  for  example,  the  taste  and  action  upon  litmus  of 
acids  and  bases  seem  to  be  due  to  the  H*  and  OH'  which  are 
the  constant  constituents  of  these  compounds,  respectively. 
Dilute  sulphuric  acid  contains  H'  and  804"  and  has  among  its 
characteristic  properties  the  evolution  of  hydrogen  when 
brought  in  contact  with  zinc.  The  concentrated  sulphuric 


INTRODUCTION  1  9 

acid  is  practically  not  ionized  and  does  not  thus  react  with 
zinc. 

The  result  of  bringing  together  two  solutions  of  ionized 
molecules  will  depend  upon  the  nature  of  the  substances  which 
can  be  formed  by  rearrangement  of  the  ions.  Anything  which 
permanently  removes  ions  from  the  field  tends  to  permanent 
chemical  change.  Thus  in  solutions  of  KOH  and  HC1  we 
have  K*  +  OH'  and  H*  +  Cl'.  It  is  possible  for  the  H*  to 
unite  with  OH'  to  form  water  and,  as  the  water  is  but  ex- 
tremely slightly  ionizable,  the  change  must  be  a  perma- 
nent one  —  these  substances  cannot  return  to  the  ionic  form. 
The  same  thing  is  true  if  they  can  unite  to  form  a  gas  or  an 
insoluble  compound,  since  in  this  way  the  ionic  materials 
leave  the  solution,  that  is,  cannot  be  ionized  again. 

Cases  of  partial  removal  of  ions  by  conversion  into  insoluble 
compounds  are  numerous.  For  example,  in  solutions  of  lead 

nitrate  and  hydrogen  sulphid  we  find  the  ions  Pb  +  2N03' 
and  2H*  +  S".  But  since  PbS  is  an  insoluble  compound  it  is 
permanently  formed  as  soon  as  the  ions  can  react.  If,  how- 
ever, we  increase  the  nitric  acid,  precipitation  is  incomplete 
and  can  be  prevented,  or  even  the  reaction  be  reversed,  the 
PbS  dissolving  when  sufficient  H'  is  present.  Such  a  reaction 
is  reversible  and  can  be  represented  by 


Pb  +  2N03'  +  2H'  +  S"*=5PbS  +  2N03'  +  2H', 

indicating  that  the  change  can  proceed  in  either  direction,  ac- 
cording to  the  relative  amounts  of  H'  or  S"  present.  This  is 
an  illustration  of  "mass  action/'  the  effect  of  increasing  the 
mass,  or  amount,  of  one  'or  other  of  the  reacting  substances. 
According  to  the  law  of  mass  action,  the  chemical  action  of  two 
substances  upon  each  other  is  proportional  to  the  active  mass 
of  each.  When  no  excessive  quantity  of  either  is  present  an 
equilibrium  may  be  attained  and  neither  change  be  complete. 
It  has  been  experimentally  discovered  that  in  a  saturated 


20  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

solution  containing  an  ionized  compound  the  product  of  the 
concentration  of  the  ions  divided  by  the  concentration  of  the 
non-ionized  molecules  is  a  constant  quantity.  This  can  be 
represented,  for  instance,  in  the  case  of  sodium  chlorid,  by 

concentration  of  Na*  X  concentration  of  Cl7 

„  —-. -T^T-^T-  ~  =  a  constant. 

Concentration  of  NaCl 

Experiment  shows  also  that  the  ionization  of  a  compound 
is  lessened  if  to  its  solution  is  added  either  ion  of  the  com- 
pound. For  example,  the  ionization  of  sodium  chlorid  is 
decreased  if  hydrochloric  acid  is  added,  the  Cl7  of  the  HC1 
being  the  effective  agent.  If  the  concentration  of  the  origi- 
nal NaCl  solution  was  sufficiently  great,  so  that  the  union  of 
Na*  and  Cl7  to  NaCl  produced  more  of  the  salt  than  the  water 
can  hold  in  solution  a  precipitate  of  NaCl  will  result. 

Or,  representing  this  phenomenon  in  another  way  by  ref- 
erence to  the  above  fraction,  if  the  Cl7  in  the  numerator  in- 
creases (since  the  value  of  the  fraction  remains  constant), 
either  the  concentration  of  the  Na'  must  correspondingly 
decrease  or  the  denominator,  which  represents  the  concen- 
tration of  the  non-ionized  salt,  must  increase.  If  this 
latter  change  produces  more  salt  than  can  be  dissolved 
there  is  an  immediate  precipitation.  On  the  other  hand, 
a  decrease  of  Na'  can  only  be  attained  by  union  with  the 
chlorin  of  the  added  HC1;  but  the  NaCl  thus  formed  would 
result  in  the  same  precipitation  as  before. 

The  product  of  concentrations  of  ions  spoken  of  above 
is  called  the  solubility  product. 

Hydrolysis 

As  a  rule,  when  chemical  change  takes  place  in  an  aqueous 
solution,  the  water  is  but  very  slightly  ionized.  However,  if 
there  is  present  a  salt  formed  from  a  weak  acid  and  a  strong 
base  or  the  opposite,  the  ions  from  the  water  may  unite  with 
the  anion  or  the  cation  of  the  salt.  For  example,  ferric  chlorid, 
FeCl3,  turns  litmus  red  because  the  Fe*  *  *  unites  with  the  OH7 
of  the  slightly  ionized  water  to  form  Fe(OH)3  leaving  Cl7  with 
H.  from  the  water  and  the  H*  gives  the  acid  reaction.  Again, 


INTRODUCTION  21 

Na2CO3  in  solution  turns  litmus  blue;  the  hydrogen  ion  of  the 
water  unites  with  the  C03"  to  form  HCO3'  thus  leaving  OH' 
which  reacts  alkaline  (page  18).  Such  a  change  involving 
a  decomposition  of  water  is  termed  hydrolysis. 

Oxidation  and  Reduction 

Beside  those  changes  where  oxygen  is  removed  from  or 
added  to  a  substance,  chemists  sometimes  speak  of  reduction 
or  oxidation  where  other  negative  elements,  like  chlorin,  are 
so  removed  or  added,  thereby  decreasing  or  increasing  the 
valence  of  the  positive  element.  With  some  of  these  the 
reaction  occurs  when  the  dry  substances  are  heated  or 
triturated  together.  With  others  the  change  is  effected  in 
solution,  either  by  the  aid  of  heat  or  at  the  ordinary  tempera- 
ture. An  increase  of  positive  valence  in  solution  means  an 
increase  in  the  positive  charge  of  the  ion ;  hence  oxidation  may 
be  denned  as  an  increase  in  the  positive  charge  and  reduction 
as  a  decrease  in  the  charge.  It  may  take  place  suddenly  or 
may  require  a  long  time  for  its  completion.  When  the  oxidiz- 
ing agent  is  in  the  solid  form  the  union  may  go  on  so  rapidly  as 
to  cause  a  dangerous  explosion.  The  production  of  either  an 
oxidation  or  a  reduction  when  medical  substances  are  pre- 
scribed together  is  one  form  of  incompatibility. 

Incompatibility 

An  incompatibility  in  a  prescription  is  caused  by  such  a 
selection  or  combination  of  the  components  that  the  usual 
action  of  these  is  modified  or  prevented  and  undesirable  results 
follow.  The  incompatibility  may  be  a  pharmaceutical  one, 
where  there  is  a  physical  but  no  chemical  change,  such  as  the 
separation  of  soluble  matters  from  their  solution.  A  thera- 
peutical incompatibility  occurs  where  drugs  are  prescribed  to- 
gether which .  are  antagonistic  in  their  physiological  action 
upon  the  system.  The  third  class,  or  chemical  incompatibility, 


22  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

is  the  only  one  with  which  this  work  will  deal.  In  this  the  in- 
gredients of  the  prescription  act  upon  each  other  in  such  way 
as  to  produce  a  chemical  change.  It  may  not  be  a  visible  one, 
both  or  all  resulting  compounds  remaining  in  solution ;  or,  as 
is  true  in  most  instances,  it  may  be  perceptible.  Some  of  the 
most  common  examples  are  those  produced  by  the  action  oi 
oxidizing  agents  which  act  both  in  dry  mixtures,  when  ex- 
plosions often  occur,  and  in  liquids.  They  are  also  caused  by 
reducing  agents,  and  in  addition  by  those  which  without 
oxidation  or  reduction  form  insoluble  substances  or  gases. 
In  this  manner  the  nature  of  the  mixture  may  be  entirely 
changed,  perhaps  being  rendered  inert  on  account  of  lessened 
absorption  through  its  insolubility  or,  on  the  other  hand,  made 
more  active  or  even  poisonous  through  the  formation  of  new 
compounds  or  by  collecting  the  active  principles  in  a  pre- 
cipitate which  is  taken  with  the  last  doses  in  the  bottle.  At 
times  the  incompatibility  in  a  prescription  is  intentional, 
one  of  the  ingredients  being  designed  to  precipitate  another, 
such  as  the  combination  of  lime-water  and  mercuric  chlo- 
rid  in  the  preparation  of  yellow  wash.  This,  however,  is 
infrequent.  Many,  although  of  course  not  all,  of  the  chemical 
incompatibilities  of  medical  compounds  are  illustrated  by 
the  reactions  performed  in  the  course  of  qualitative  analysis. 
A  thorough  study  should  be  made  of  these  for  the  purpose 
of^discovering  such  as  are  common  or  possible. 

Where  one  or  the  other  of  the  compounds  entering  into  a 
reaction  is  a  poison,  the  results  obtained  in  the  following  tests 
may  often  be  used  in  the  selection  of  an  antidote.  This  may 
act  by  forming  an  insoluble  compound  with  the  poison  and 
thereby  preventing  its  absorption,  by  forming  an  inert  com- 
pound through  union,  or  by  more  complete  decomposition  of 
the  .poisonous  substance.  Only  a  knowledge  of  their  chemi- 
cal properties  will  give  the  physician  a  mastery  over  the 
subject  of  poisons. 


PART  I 

QUALITATIVE  ANALYSIS 


CHAPTER  I 

METALS  (CATIONS) 

THE  PREPARATION  OF  SOLUTIONS  FOR  ANALYSIS 

IN  dissolving  solids  for  analysis  no  general  rule  can  be  given 
except  that  the  weakest  possible  solvent  should  be  employed, 
and  of  that  only  as  much  as  is  necessary.  Heat  is  generally 
of  great  assistance  in  effecting  solution.  It  is  advisable  first 
to  learn  the  best  solvent  by  testing  small  amounts  of  the  solid, 
and,  when  this  has  been  ascertained,  to  use  as  much  as  is 
desirable  for  the  analysis.  Solvents  may  be  tried  in  a  test- 
tube  in  the  following  order,  always  warming  the  liquid,  and 
using  separate  portions  of  solid  for  each  test. 

1.  Water. 

2.  Dilute  hydrochloric  acid. 

3.  Concentrated    hydrochloric    acid    (2-4    c.c.).     Warm 
gently,  afterward  dilute  with  water  and  boil. 

4.  Nitric  acid,  first  dilute,  then  concentrated. 

5 .  The  residue  from  3  in  dilute,  then  concentrated  nitric  acid . 

6.  Aqua  regia. 

If  concentrated  acid  is  necessary  it  should  be  largely  re- 
moved, by  evaporating  nearly  to  dryness  under  a  hood,  before 
proceeding  to  the  analysis.  If  the  solid  dissolves  in  two  sepa- 
rate solvents,  the  solutions  may  be  mixed  for  the  analysis, 
provided  this  does  not  cause  precipitation;  or  the  portions 
may  be  tested  separately. 

23 


24  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

Some  substances  which  remain  insoluble  after  the  above 
treatment  are  the  sulphates  of  barium,  strontium,  lead,  and 
possibly  calcium;  the  chlorid,  bromid,  iodid,  and  cyanid  of 
silver;  the  oxids  of  tin,  aluminum,  chromium,  silicon,  and 
possibly  iron;  calcium  fluorid;  some  silicates;  carbon  and 
sulphur.  The  residue  can  be  fused  on  platinum  with  three 
times  as  much  dry  sodium  carbonate,  the  mass  boiled  in 
water,  filtered,  and  the  filtrate  tested  for  anions.  Wash  the 
residue,  dissolve  in  dilute  nitric  acid  and  test  for  cations. 

GROUP  V 

The  Metals  of  the  Alkalies,  Potassium,  Sodium  (and 
Ammonium) 

Almost  all  compounds  of  this  group  of  metals  are  soluble 
in  water,  consequently  they  can  rarely  be  identified  by  the 
formation  of  precipitates.  Their  hydroxids,  carbonates,  and 
also  sulphids  when  dissolved  in  water  are  alkaline  in  reaction. 

Potassium,  K 

Use  a  5-per  cent,  solution  of  KC1  for  the  following  reactions. 

1.  Dip  a  looped  platinum  wire  into  a  potassium  solution 
and  hold  it  in  the  outer  part  of  a  Bunsen  flame.     The  color  of 
the  flame  above  the  substance  becomes  bluish-violet.    Look 
at  the  flame  through  one  or  more  thicknesses  of  blue  glas' 
(cobalt  glass).     The  color  is  not  destroyed  by  the  glass,  but 
becomes  a  more  distinct  violet.     No  other  metal  will  give  these 
results,  although  many  organic  compounds  give  a  luminous 
flame  which  appears  of  a  similar  color  through  the  blue  glass. 
The  organic  matter  may  be  first  destroyed  by  burning  and 
then  the  residue  tested  for  potassium. 

2.  When  a  solution  of  sodium  cobaltinitrite  is  added  to  one 
of  potassium,  the  potassium  ion  unites  with  the  cobaltinitrite 
ion,  Co(NO2)6'".,  to  form  a  yellow  precipitate  of  potassium 


METALS  (CATIONS)  25 

cobaltini trite,  K3Co(N02)6,  which  is  somewhat  soluble  in 
water,  and  therefore  may  not  appear  except  from  rather  con- 
centrated solutions. 

3.  Platinic  chlorid,  to  which  hydrochloric  acid  has  been 
added,  contains  the  chlorplatinate  ion,  PtCV  and  precipi- 
tates the  potassium  ion  from  neutral  or  alkaline  solutions,  if 
they  are  not  too  dilute,  in  yellow  octahedral  crystals  of 
potassium  chlorplatinate,  K2PtCl6.     The  addition  of  alcohol 
renders  the  precipitation  more  complete. 

Sodium,  Na 

Use  a  5-per-cent.  solution  of  NaCl  for  the  test. 

4.  Compounds  of  sodium  give  an  intense  yellow  color  to 
the  flame  when  volatilized  in  it,  as  in  the  potassium  tests,  on  a 
platinum  wire.     This  color  does  not  pass  through  a  blue 
glass  providing  the  latter  is  of  sufficient  thickness. 

Ammonium,  NH4 

Use  a  5-per-cent.  solution  of  NH4C1  for  the  reactions. 

5.  When  ammonium  salts,  either  solid  or  in  solution,  are 
heated  with  sodium  hydroxid,  NaOH,  they  are  decomposed, 
ammonia,  NH3,  being  liberated. 

NH4C1  +  NaOH  =  NH3  +  NaCl  +  H2O. 

This  is  identified  by  its  characteristic  odor  and  also  by  its 
turning  blue  a  moistened  piece  of  red  litmus-paper  suspended 
in  the  mouth  of  the  test-tube  in  such  a  manner  that  it  does 
not  touch  the  inside  of  the  tube.  The  hydroxids  of  the 
alkaline  earths  and  also  of  potassium  will  likewise  set  am- 
monia free  from  its  compounds. 

The  explanation  of  the  evolution  of  NH3  from  solutions  of 
the  ammonium  ion  is  that  NaOH  being  a  strong  base  is,  in 
solution,  largely  ionized  to  Na'  and  OH'.  Since  NH4OH  is  a 
weak  base,  it  ionizes  but  slightly  and  consequently  forms 


26  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

when  NH4.  and  OH'  come  together.     Being  very  unstable  at  a 
high  temperature,  it  is  decomposed  by  heating. 

NH4OH  =  NH3  +  H20. 

6.  The  ammonium  ion,  like  the  potassium  ion,  forms  a 
yellow  precipitate  (NH4)3Co(NO2)6,  with  the  cobaltini trite 
ion.     It  does  not  appear  in  very  dilute  solutions  because  of 
its  solubility  in  water. 

7.  A  platinic  chlorid  solution,  containing  the  chlorplatinate 
ion,  PtCl6",  precipitates  from  neutral  or  acid  solutions  of  am- 
monium salts,  when  they  are  sufficiently  concentrated,  yellow 
octahedral   crystals   of   ammonium   chlorplatinate    (NH4)2- 
PtCl6.       It    is    similar    to    the    corresponding    potassium 
compound. 

8.  All  dry  ammonium  compounds  are  volatilized  or  decom- 
posed when  heated  on  platinum  foil  or  a  piece  of  porcelain. 

9.  A  few  drops  of  Nessler's  reagent  added  to  10  c.c.  of  a 
solution  of  an  ammonia  compound  gives  a  brown  precipitate, 
NHg2I.     The  test  is  so  sensitive  that  one  part  of  ammonia 
can  be  detected  in  a  million  of  water.     With  such  dilute 
solutions  there  is  no  precipitate,  but  only  a  yellow  to  brown 
color  produced. 

Practical  Exercise  in  the  Analysis  of  Group  V 

Mix  in  a  test-tube  5  c.c.1  each  of  solutions  of  potassium, 
sodium  and  ammonium.  Warm  half  of  this  with  sodium 
hydroxid  and  observe  that  ammonia  is  set  free  as  from  the 
simple  solution.  Make  the  flame  test  with  the  other  half, 
observing  the  color  both  with  and  without  the  blue  glass. 
The  sodium  yellow  is  seen  without  the  glass  and  the  potas- 
sium violet  with  it,  neither  interfering  with  the  other. 

1  It  is  advisable  that  the  student  should,  at  the  beginning  of  his  course, 
determine  the  volumes  of  his  test-tubes  and  other  vessels  in  metric  measures 
and  thereafter  use  only  these  measures  in  chemical  work. 


METALS  (CATIONS)  27 

Questions  for  Further  Study  on  Group  V 

To  be  answered  by  the  student. 

How  do  you  explain  the  destruction  of  the  yellow  light  by 
the  cobalt  glass,  while  the  violet  passes  through  unchanged? 
What  kind  of  chemical  reagents  turned  red  litmus  blue  ?  Why 
does  the  paper  act  more  quickly  if  moist  ?  Why  should  it  not 
be  allowed  to  touch  the  inner  wall  of  the  test-tube?  Could  the 
practical  exercise  above  be  used  to  determine  the  composition 
of  an  unknown  mixture  where  K',  Na*  and  NH4'  might  be  pre- 
sent or  absent  ?  Can  the  dry  sodium  and  potassium  compound 
be  freed  from  ammonium  compounds  by  heating  to  a  high  tem- 
perature ?  What  compounds  of  the  alkali  metals  would  be  in- 
compatible with  acids  ?  What  would  be  the  best  for  neutraliz- 
ing acids  ?  What  property  of  a  chemical  compound  prevents 
its  appearance  as  a  precipitate?  What  compounds  of  am- 
monium can  be  identified  by  the  odor?  Which  of  the  com- 
pounds of  the  alkali  metals  are  suitable  for  internal  use  as 
antidotes  in  case  of  poisoning  by  the  mineral  acids? 

GROUP  IV 

The  Metals  of  the  Alkaline  Earths,  Magnesium,  Calcium, 
Strontium  and  Barium 

Their  oxids  are  soluble  in  water  (Mg  only  slight  so),  and 
the  solution  has  an  alkaline  reaction.  The  phosphates  and 
carbonates  are  insoluble  in  water,  but  soluble  in  acids,  except 
those  of  barium,  strontium,  and  calcium  in  sulphuric,  the  car- 
bonate yielding  CC>2  and  neutralizing  the  acid.  None  of 
them  form  insoluble  compounds  by  the  action  of  hydrogen 
sulphid  or  ammonium  sulphid. 

Take  for  each  test  3-4  ex.  of  the  solution  of  their  salts. 

Magnesium,  Mg 

Use  a  5-per-cent.  solution  of  MgCl2  or  MgSO4  for  the 
reactions. 


28  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

10.  The  hydroxid  ion  (for  instance  from  solutions  of  NaOH, 
KOH,  NH4OH,  or  Ba(OH)2)  forms,  with  the  magnesium  ion, 
magnesium  hydroxid,  Mg(OH)2,  a  white  gelatinous  precipitate. 

Mg  +  2OH'  =  Mg(OH)2,  or 
MgCl2  +  2NaOH  =  Mg(OH)2  +  2NaCl. 

If  NH4OH  is  the  reagent  the  precipitation  is  incomplete 
because  complex  soluble  ammonium-magnesium  ions  are 
formed. 

This  precipitation  does  not  occur  if  to  the  solution  has  been 
added  an  ammonium  salt  since,  as  has  been  shown  (page  20), 
the  common  ion,  NH4,  so  far  decreases  the  ionization  of  the 
NH4OH  that  the  product  of  the  concentration  of  the  magne- 
sium ion  and  the  hydroxid  ion  is  less  than  the  solubility  value. 
In  all  such  cases  no  precipitate  forms. 

1 1 .  The  carbonate  ion,  found  conveniently  in  solutions  of 
ammonium   carbonate    (NH4)2CO3,    or   sodium   carbonate, 
Na2COs,  produces  basic  carbonates  varying  in  composition, 
types  of  which  are  MgCO3,  Mg(OH)2,  and  (NH4)2Mg(CO3)2. 
Like  the  magnesium  hydroxid,  these  carbonates  are  soluble  in 
solutions  of  ammonium  salts.     They  will  therefore  disappear 
upon  the  addition  of  ammonium  chlorid  to  the  liquid  which 
contains  them  after  heating,  or,  if  ammonium  chlorid  is  added 
to  the  magnesium  solution  before  the  carbonate  ion,  there  is 
no  precipitation. 

These  precipitates  are  very  light,  becoming  heavier  when 
heated  in  the  liquid  in  which  they  are  formed. 

12.  The  phosphate  or  hydrophosphate  ion,  HPO4",  con- 
tained in  a  solution  of  sodium  phosphate,  Na2HPO4  converts 
the  magnesium  ion  into  insoluble,  white  MgHPO4.     If  am- 
monium hydroxid  is  present  also  the  ammonium  ion  unites  to 
form  white  crystalline  ammonium  magnesium  phosphate, 
NH4MgP04.     Some  ammonium  salt,  like  ammonium  chlorid, 
should  be  added  in  order  to  prevent  the  union  of  the  mag- 


METALS  (CATIONS)  29 

nesium  and  hydroxid  ions  and  the  consequent  formation  of 
magnesium  hydroxid  instead  of  the  phosphate  (10). 

Mg4-HP04//+NH4/+OH/  =  NH4MgP04+H20  or,  repre- 
senting all  the  original  compounds  and  the  products  which 
might  be  obtained  by  evaporation  from  the  solution, 
MgCl2+Na2HPO4+NH4OH  =  NH4MgPO4+ 2NaCl+H2O 

Under  the  microscope  the  crystals  are  seen  to  have  a 
stellate  or  fern-leaf  form,  especially  if  quickly  precipitated. 
When  formed  slowly  they  are  prismatic.  All  precipitates 
of  magnesium  are  soluble  in  acids. 

Calcium,  Ca 

Use  for  reactions  a  5-per-cent.  solution  of  CaCl2. 

13.  Calcium  salts  when  held  in  the  blue  Bunsen  flame  give 
it  a  yellowish-red  color.     This  is  destroyed  by  the  blue  glass, 
through  a  thin  piece  of  which  it  appears  grayish-green.     The 
flame  test  is  most  marked  with  the  chlorid  or,  when  other 
salts  are  tested,  after  moistening  with  hydrochloric  acid. 

14.  The  carbonate  ion,  from  such  solutions  as  ammonium 
carbonate,  precipitates  the  calcium  ion  as  calcium  carbonate, 
CaCOs,  white  and  amorphous,  and  practically  insoluble  in 
ammonium  salts.     If,  after  settling,  the  solution  is  separated 
from  the  precipitate,  the  latter  will  dissolve  with  efferves- 
cence when  an  acid  is  poured  on  it. 

15.  The  oxalate  ion  of  the  ammonium  oxalate  solution 
produces  a  fine,  white  precipitate  of  calcium  oxalate,  CaC2O4 
from  neutral  or  alkaline  calcium  solutions. 

Ca+C204"  =  CaC204 

or,  representing  all  the  elements  present  except  the  solvent, 
water, 


30  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

It  is  soluble  in  hydrochloric  or  nitric  acid. 

1 6.  The  sulphate  ion,  from  a  solution  of  dilute  sulphuric 
acid  or  any  sulphate  except  that  of  calcium,'  precipitates  from 
a  concentrated  solution,  fine,  white,  calcium  sulphate,  CaSO4, 
which  is  soluble  in  a  large  amount  of  water.     It  is,  in  conse- 
quence, not  precipitated  from  dilute  solutions.     The  calcium 
ion  cannot  be  precipitated  by  a  solution  of  calcium  sulphate. 

17.  The  phosphate  ion  throws  down  from  calcium  solutions 
a  white,  flocculent  acid  calcium  phosphate,  CaHP04. 

18.  The  dichromate  ion  does  not  precipitate  the  calcium 
ion. 

Strontium,  Sr 

A  5-per-cent.  solution  of  SrCl2  may  be  used  for  the  reactions. 

19.  Strontium  salts  give  a  deep, crimson  color  to  the  blue 
flame,  best  after  the  addition  of  hydrochloric  acid.     This  is 
seen  through  a  thin  blue  glass  but  not  through  a  thick  one. 

20.  The  carbonate  ion  gives  a  white,  flocculent  precipitate 
of  SrCOs,  similar  in  properties  to  CaCOs. 

2 1 .  The  oxalate  ion  precipitates  strontium  oxalate,  SrC2O4, 
if  the  solution  is  not  very  dilute.     It  is  a  fine  white  powder. 

2  2 .  The  sulphate  ion  precipitates  strontium  as  the  sulphate, 
SrS04.  The  strontium  sulphate  is  somewhat  soluble  in  water, 
consequently  it  is  not  completely  precipitated  and  appears 
rather  slowly  in  dilute  solutions.  It  is  less  soluble  in  alcohol. 

23.  The  phosphate  ion  produces  a  white,  flocculent  precipi- 
tate of  SrHPO4,  acid  strontium  phosphate. 

24.  The  dichromate  ion  does  not  precipitate  strontium 
from  its  neutral  or  acid  solutions. 

Barium,  Ba 

All  soluble  barium  compounds  act  as  poisons. 

Use  for  the  reactions   a   5-per-cent.   solution  of  BaCl2. 

25.  Barium  chlorid  and  most  other  barium  salts  to  which 


METALS  (CATIONS)  31 

hydrochloric  acid  has  been  added  produce  a  yellowish-green 
flame  when  held  in  the  oxidizing  flame  of  a  Bunsen  burner. 

26.  The  carbonate  ion  precipitates  barium  from  its  solution 
as  BaC03,  similar  in  its  properties  to  the  carbonates  of  cal- 
cium and  strontium. 

27.  The  oxalate  ion  forms  barium  oxalate,  BaC2O4,  a  fine, 
white,  heavy  precipitate. 

28.  The  sulphate  ion  precipitates  immediately  the  barium 
from  its  solution  as  a  heavy,  very  fine,  white  solid — barium 
sulphate,  BaSO4. 

29.  The  phosphate  ion  yields  barium  phosphate,  BaHPO4, 
a  compound  similar  in  composition,  formation  and  properties 
to  the  corresponding  salts  of  calcium  and  strontium. 

30.  The  dichromate  ion  or  the  chromate  ion  from  the  solu- 
tions of  potassium  dichromate  (bichromate),  or  potassium 
chromate    precipitates    bright    yellow    barium    chromate, 
BaCr04,  insoluble  in  acetic  acid. 

Directions  for  the  Separation  of  the  Cations  of  Groups 

IV  and  V 

If  the  solution  which  contains  the  ions  of  groups  IV  and  V 
is  not  already  alkaline  make  it  so  by  adding  a  little  ammo- 
nium hydroxid,  then  about  5  c.c.  of  ammonium  chlorid  and, 
finally,  ammonium  carbonate  as  long  as  a  precipitate  is 
formed.  Warm  nearly  to  the  boiling-point  and  filter.  Wash 
the  precipitate,  discarding  the  wash-water. 

The  precipitate  contains  the  carbonate  of  barium,  stron- 
tium, and  calcium.  The  filtrate  contains  magnesium,  potas- 
sium, and  sodium. 

Dissolve  the  precipitate  on  the  paper  by  pouring  over  it 
5-10  c.c.  of  hot  acetic  acid,  using  the  same^acid  repeatedly  if 
the  first  application  is  insufficient.  To  the  resulting  solution 
add  potassium  dichromate  which  precipitates  the  barium  as 


32  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

yellow  barium  chromate.  Filter  and,  after  diluting  the  ni- 
trate with  water  to  about  25  c.c.,  add  dilute  sulphuric  acid 
to  precipitate  the  strontium.  The  precipitate  of  strontium 
sulphate  is  a  very  fine  white  solid.  Let  it  settle  five  minutes. 
Filter,  wash,  moisten  the  precipitate  on  the  filter  with  a  few 
drops  of  hydrochloric  acid  and  confirm  the  presence  of  stron- 
tium by  the  deep  crimson  color  imparted  to  the  flame  of  a 
Bunsen  burner  when  the  substance  is  held  in  it  on  a  platinum 
wire.  It  is  also  reddish  through  thin  cobalt  glass.  Make 
the  filtrate  from  the  strontium  alkaline  with  ammonium  hy- 
droxid,  then  add  ammonium  oxalate.  Calcium  is  precip- 
tiated  as  fine,  white  calcium  oxalate.  It  may  be  confirmed 
by  the  reddish-yellow  color  of  its  flame,  testing  in  the  same 
manner  as  for  strontium.  It  appears  a  dirty  green  through 
a  thin  blue  glass. 

The  solution  containing  magnesium,  potassium,  and  so- 
dium should  be  tested  first  for  the  sodium  and  potassium  by 
the  color  imparted  to  a  Bunsen-burner  flame — yellow  from 
sodium,  and  violet  when  potassium  is  present  and  the  flame 
is  viewed  through  a  sufficiently  thick  blue  glass.  The  test  is 
more  delicate,  with  small  amounts  of  these  metals,  if  the  liquid 
is  first  concentrated  to  a  few  cubic  centimeters  by  boiling. 

To  detect  magnesium  add  to  the  solution  which  has  been 
tested  for  sodium  and  potassium  a  little  sodium  phosphate. 
Shake  vigorously  and  let  it  stand  without  warming.  Mag- 
nesium ammonium  phosphate  is  precipitated — small,  white 
snowflake-shaped  crystals  when  seen  with  the  microscope. 
An  excess  of  ammonium  hydroxid  favors  the  precipitation. 

The  presence  of  ammonium  in  the  original  substance  can 
be  determined  only  by  applying  the  test  to  this  directly. 
This  is  done  by  adding  5  c.c.  of  sodium  hydroxid,  or  enough 
to  give  an  alkalin^reaction,  and  boiling.  The  ammonia  gas 
which  is  evolved  can  be  identified  by  its  characteristic  odor 
or  by  its  turning  red  litmus-paper  blue. 


r 


METALS  (CATIONS) 


33 


TABLE  I 

OUTLINE  OF  SEPARATION  OF  CATIONS  OF  GROUPS  IV  AND  V    • 

Heat  a  small  portion  of  the  original  substance  with  sodium  hydroxid. 
An  evolution  of  ammonia  gas  indicates  ammonium. 

Make  the  remainder  of  the  solution  which  contains  groups  IV  and  V 
alkaline  with  ammonium  hydroxid  and  add  ammonium  chlorid  and  ammo- 

_.- u___._     _    «...     .!,_..,  .          ^ 


nium  carbonate,  heat,  filter,  and  wash. 


Precipitate  contains  BaCO3,  SrCO3,  CaCO3. 
Dissolve  in  dilute  acetic  acid  and  add  potassium 
dichromate.     Filter. 

Filtrate  contains  Mg,   K 
and    Na.     Test    color    of 
flame.        Yellow   indicates 
Na.     Violet    through    the 
blue     glass     indicates     K. 
Add  sodium  phosphate.    A 
white,    crystalline   precipi- 
tate indicates  Mg. 

A  yellow 
precipitate 
is  BaCrO4. 

Solution  contains  Sr  and  Ca.     Add 
dilute  H2SO4. 

A    fine    white 
precipitate  is 
SrSO4.      Confirm 
by  crimson  flame. 

Solution     con 
tains  Ca.     Make 
alkaline  with  am- 
monium hydroxid 
and  add  ammo- 
nium oxalate.    A 
fine  white  precipi- 
tate   is    CaC2O4. 
Confirm  by  red- 
dish-yellow flame. 

Explanations  of  the  Operations  Used  in  the  Separation 
of  the  Cations  of  Groups  IV  and  V 

The  ammonium  chlorid  must  be  used  here  to  prevent  the 
precipitation  of  magnesium  with  the  others  of  the  alkaline 
earth  metals  (10) .  Ammonium  hydroxid  is  necessary  to  neu- 
tralize any  acid  that  may  be  present  as  this  would  decompose 
the  ammonium  carbonate  and  prevent  the  precipitation  of 
any  of  the  ions.  The  carbonates  of  barium,  strontium,  and 
calcium  which  are  formed  in  cold  solutions  are  not  completely 
precipitated  if  carbon  dioxid  is  present  because  of  the  forma- 
tion of  soluble  acid  carbonates,  but  these  are  converted  by 
heating  into  insoluble  ones. 

The  acetic  acid  displaces  the  carbonid>^id  from  the  pre- 
cipitated carbonates,  the  positive  part  of  the  compounds 
going  into  solution  as  cations.  The  majority  of  precipitates 
are  not  soluble  with  sufficient  ease  to  allow  this  method  of 

3 


34  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

solution,  but  must  be  removed  from  the  filter  and  mixed,  or 
often  heated,  with  the  solvent.  The  effervescence  is  caused 
by  the  evolution  of  C02  which  is  always  set  free  by  the  action 
of  an  acid  on  a  carbonate.  If  no  water  were  added  to  the 
solution  of  strontium  and  calcium  before  the  former  was  pre- 
cipitated as  a  sulphate  some  or  most  of  the  calcium  might  also 
be  thrown  down  (16).  Strontium  is  slowly  precipitated  on 
account  of  the  solubility  of  its  sulphate  (22),  therefore,  it  is 
necessary  to  allow  sufficient  time  for  this  to  be  accomplished. 

If  the  strontium  sulphate  were  not  washed  some  of  the  cal- 
cium solution  would  remain,  if  this  metal  were  present  in  the 
mixture,  and  the  color  of  the  flame  would  be  modified.  The 
acid  makes  it  more  distinct. 

Ammonium  oxalate  does  not  precipitate  bases  of  this  group 
from  acid  solutions,  such  as  this  one  is  after  the  use  of  sul- 
phuric acid,  consequently  the  latter  must  be  neutralized  in 
order  to  detect  the  calcium. 

The  flame  test  for  the  alkalies  should  not  be  made  until 
the  removal  of  the  barium,  strontium,  and  calcium  ions  from 
the  mixture  because  of  the  effect  they  would  give  if  they  were 
present.  The  sodium  would  not  be  concealed,  but  the  potas- 
sium might  be  difficult  of  identification  (19). 

The  test  for  sodium  must  be  applied  before  that  for  magne- 
sium because  the  precipitation  of  magnesium  requires  a  sod- 
ium salt  and  the  sodium  from  this  remains  in  solution  and 
passes  into  the  filtrate  from  the  magnesium.  With  solid 
sodium  compounds  or  very  concentrated  solutions  some  of 
the  light  may  pass  through  the  cobalt  glass  if  this  is  not  suf- 
ficiently thick,  but  it  then  appears  blue,  and  not  violet,  as 
can  be  seen  by  comparing  it  with  a  known  potassium  solution. 

The  sodium  phosphate  precipitates  the  magnesium  very 
completely,  but  if  tjiere  is  not  much  magnesium  present  a 
long  time  of  standing  may  be  necessary. 

The  original  substance  can  be  tested  for  ammonium  by 


METALS  (CATIONS)  35 

adding  sodium  hydroxid  to  it  directly  without  dissolving  as 
the  ammonia  will  be  readily  set  free.  If  only  a  minute 
amount  is  present  it  can  be  detected  by  hanging  the  moist  red 
litmus-paper,  in  such  a  manner  that  it  shall  not  touch  the 
tube,  from  a  cork  which  closes  the  test-tube  and  allowing  it 
to  stand  for  twenty-four  hours  at  the  ordinary  temperature. 
The  test  with  litmus-paper  is  much  more  sensitive  than  iden- 
tification by  the  odor.  In  this  case  care  must  be  taken  that 
the  reagent  contains  no  ammonium  compounds.  It  is  neces- 
sary to  use  the  original  substance  to  test  for  ammonia  and  not 
the  filtrate  from  barium,  strontium  and  calcium  carbonates 
where  we  might  naturally  look  for  it,  since  before  getting 
this  filtrate  a  number  of  ammonium  compounds  have  been 
added  to  the  solution. 

Practical  Exercises  in  the  Separation  of  the  Cations  of 
Groups  IV  and  V 

Mix  in  a  large  test-tube  3-4  c.c.  of  the  solution  of  each 
and  analyze  the  mixture  of  the  seven  according  to  Table  I. 

In  the  same  manner  make  analyses  of  unknown  mixtures 
of  these  groups,  which  can  be  obtained  from  the  instructors, 
and  report  results  in  writing. 

Questions  for  Further  Study  on  Group  IV 

Why  are  the  metals  of  group  IV  called  metals  of  the  alka- 
line earths  ?  When  their  oxids  dissolve  in  water  is  there  any 
chemical  change  ?  What  is  such  a  solution  of  calcium  called  ? 
For  what  is  it  used?  What  property  of  these  compounds  is 
made  use  of  to  separate  the  cations  of  this  group  from  those  of 
group  V?  What  is  meant  by  effervescence?  Why  might 
it  be  expected  that  the  flame  of  calcium  compounds  would  be 
more  plain  by  the  addition^of  hydrochloric  acid?  What 
would  be  the  effect  of  adding  alcohol  to  a  mixture  of  sulphuric 
acid  with  a  dilute  solution  of  strontium?  Why  does  not  a 


36  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

solution  of  calcium  sulphate  precipitate  concentrated  solu- 
tions of  calcium  salts  when  this  can  be  done  by  the  use  of 
other  soluble  sulphates?  Why  will  an  ammonium  oxalate 
solution  not  precipitate  calcium  from  a  strongly  acid  solu- 
tion? In  the  analysis  of  group  IV  what  would  be  indicated, 
if  instead  of  a  crystalline  precipitate  where  we  should  expect 
magnesium,  a  white  amorphous  precipitate  should  appear? 
Are  there  any  other  alkaline  gases  except  ammonia?  Which 
compounds  of  the  metals  of  group  IV  can  be  employed  for 
the  neutralization  of  acids?  Are  any  ever  used  medicinally 
for  this  purpose?  Can  an  acid  be  neutralized  by  a  substance 
which  is  itself  neutral  in  reaction?  What  use  is  made  of  cal- 
cium carbonate  in  medicine?  By  what  other  name  is  mag- 
nesium carbonate  known,  and  for  what  properties  is  it 
valuable  in  medicine?  What  is  the  difference  in  composition 
between  the  soluble  acid  carbonate  of  calcium  and  the  in- 
soluble carbonate  of  the  same  metal  ?  Where  does  the  former 
occur  naturally? 

GROUP  m 

Aluminum,  Chromium,  Zinc,  Manganese,  Cobalt,  Nickel  and  Iron 
Not  only  the  carbonates  and  phosphates  of  this  group  are 
insoluble  in  water,  but  also  the  oxids,  hydroxids,  and  sulphids, 
except  the  sulphids  and  carbonates  of  aluminum  and  chro- 
mium which  cannot  be  formed  by  precipitation  but  are 
hydrolyzed  to  the  hydroxids.  The  ions  of  the  group  are 
therefore^ precipitated  from  neutral  solutions  by  hydroxid, 
carbonate,f  phosphate,  and  sulphid  ions,  the  carbonate  and 
sulphid  ions  forming  hydroxids  of  aluminum  and  chromium 
by  hydrolysis.  Hydrogen  sulphid  does  not  precipitate  ions 
of  this  group  from  acid  solutions. 

Aluminum  (Aluminium),  Al 

Aluminum  dissolves  in  hydrochloric  acid  and  caustic 
alkalies,  but  only  slowly  in  other  acids.  Aluminum  acts  as  a 
cation,  Al'",  also  as  part  of  an  anion  like  A1(V  or  A10'". 


METALS  (CATIONS)  37 

Use  for  the  reactions  of  the  cation  a  5-per-cent.  solution  of 
A12(S04)3  or  KA1(S04)2. 

31.  The  sulphid  ion,  which  can  be  obtained  from  a  solution 
of  ammonium  sulphid,  produces  in  solutions  of  the  cation  a 
very  light,  gelatinous,  white  precipitate  of  aluminum  hy- 
droxid,  Al(OH),. 

Aluminum  sulphid,  the  formation  of  which  might  be  ex- 
pected, cannot  exist  in  water,  but  hydrolyzes  (page  20),  yield- 
ing the  hydroxid  and  hydrogen  sulphid,  A12S3  +  3H2O  = 
A1(OH)3  +  3H2S.  Hydrogen  sulphid  produces  no  precipi- 
tate with  the  aluminum  ion  unless  the  H'  is  neutralized  by 
adding  the  hydroxid  ion,  for  example,  from  NH4OHorNa- 
OH.  Aluminum  hydroxid  is  readily  soluble  in  acids. 

A12(S04)3+  3(NH4)2S  +  6H20  =  2A1(OH)3  + 
3(NH4)2S04  +  3H2S. 

32.  When  barium  carbonate,  suspended  in  water,  is  mixed 
with  an  aluminum  solution  the  liquid,  after  filtration  or  set- 
tling of  the  residue,  will  be  found  free  from  aluminum. 

33.  The  hydroxid  ion  precipitates  the  cation,  Al'",  as  the 
hydroxid,  A1(OH)3,  described  in  31.     This  dissolves  immedi- 
ately in  an  excess  of  the  reagent,  if  sodium  hydroxid  or  potas- 
sium hydroxid  has  been  used,  the  aluminum  then  being 
found  in  the  anion,  as  NaAlO2,  sodium  aluminate.     If  the 
reagent  is  ammonium  hydroxid  the  aluminum  hydroxid  is  but 
slightly  soluble  and  can  be  completely  precipitated  by  heat- 
ing.    With  sodium   hydroxid   or  potassium  hydroxid^  the 
aluminate  formed  is  not  decomposed  by  heating  and  the  hy- 
droxid does  not  precipitate  on  boiling.     It  can  be  precipitated 
by  adding  ammonium  chlorid,  best  by  the  aid  of  heat.     Acidi- 
fying the  solution  of  the  aluminate  destroys  it,  leaving  the 
aluminum  as  the  cation. 

34.  The  carbonate  ion  with  solutions  of  the  aluminum  ion 
(cation)  likewise  forms  aluminum  hydroxid,  hydrolysis  oc- 
curring as  with  the  sulphid,  and  C02  escaping, 


38  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

2AL-  +  6HCCV  =  2A1(OH)3  +  6C02j  or 
2KA1(S04)2  +  6NaHC03  =  2A1(OH)3  + 
K2SO4  +  3Na2S04  +  6C02. 

Chromium,  Cr 

Chromium  may  be  met  with  in  two  classes  of  compounds 
the  salts,  where  it  occupies  the.positive  part  of  the  compound, 
forming  the  cation  in  solutions,  such  as  CrCl3  or  Cr2(SO4)3; 
and  the  chromates,  or  chromic  acid  derivatives,  where  it.  is 
found  in  the  negative  part,  such  as  PbCr04  or  K2Cr2O7. 
With  these  two  classes  most  reagents  give  different  results. 
Only  the  former  will  be  considered  in  this  place. 

The  chromic  salts  have  a  green  or  violet  color,  the  chromic 
ion  is  also  green. 

Use  for  the  reactions  of  the  chromic  salts  a  5-per-cent.  solu- 
tion of  KCr(S04)2. 

35.  The  sulphid  ion  (from  ammonium  sulphid)  precipitates 
the  chromic  ion  as  chromium  hydroxid,  Cr(OH)3,  a  grayish- 
green  or  bluish-green,  voluminous  gelatinous  compound.     It 
is  soluble  in  acids.     Hydrogen  sulphid  is  formed  by  hydrolysis 
as  it  is  by  the  action  of  ammonium  sulphid  on  aluminum  salts. 

36.  The  hydroxid  ion  precipitates  the  chromic  ion  as  chro- 
mium hydroxid.     It  is  soluble  in  excess  of  the  reagent,  impart- 
ing a  pink  or  green  color  to  the  liquid,  but  is  precipitated 
from  this  solution  as  chromic  hydroxid  by  boiling. 

37.  Dry  compounds  of  chromium  if  mixed  with  dry  potas- 
sium or  sodium  nitrate  and  sodium  carbonate  and  then  fused 
on  the  platinum  foil  are  thereby  converted  to  yellow  potas- 
sium  or   sodium   chromates.     These   are   soluble  in  water, 
forming  a  bright  yellow  liquid  and  giving  the  reactions  of 
chromic  acid.     Thus,  after  acidifying  with  acetic  acid,  a  drop 
of  lead  acetate  solution  gives  a  yellow  precipitate  of  lead 
chromate,  PbCr04.     For  this  test  any  of  the  precipitates  may 
be  employed  after  filtering  from  the  solution  and  drying  on 
the  foil 


METALS  (CATIONS)  39 


2Cr(OH)3  +  3NaNO3  +  2Na2CO3  =  2Na2Cr0.i  +  3 
2CO2  +  3NaN02. 

38.  Barium  carbonate  precipitates  basic  chromium  salts 
in  the  same  manner  as  it  does  those  of  aluminum. 

Zinc,  Zn 

Zinc  is  readily  soluble  in  hydrochloric,  nitric,  and  sulphuric 
acids.  Its  salts  are  colorless. 

For  the  reactions  a  2-per-cent  solution  of  ZnS04  may  be 
used. 

39.  The  sulphid  ion,  obtained  best  from  ammonium  sul- 
phid,  precipitates  the  zinc  ion,  from  its  neutral  or  akaline 
solutions  as  the  sulphid,  ZnS.  It  is  a  white  flocculent  solid, 
insoluble  in  excess  of  the  reagent  or  in  alkalies,  but  soluble  in 
most  mineral  acids.  With  hydrogen  sulphid  the  precipita- 
tion is  incomplete. 

40.  The  hydroxid  ion  precipitates  zinc  hydroxid,  Zn(OH)2, 
a  white  gelatinous  compound  readily  soluble  in  excess  of 
ammonia.     On  boiling  the  zinc  hydroxid  precipitates  again  in 
whole  or  in  part  unless  ammonium  salts  are  present,  which 
will   prevent   the   separation.     From   solutions   in   sodium 
hydroxid  or  potassium  hydroxid  it  does  not  precipitate  on 
boiling,  if  a  large  excess  of  the  alkali  is  present. 

41.  The  ferrocyanid  ion  (potassium  ferrocyanid)  precipi- 
tates zinc  ferrocyanid,  Zn2Fe(CN)6,  a  white  compound,  in- 
soluble in  cold  dilute  hydrochloric  acid  but  somewhat  soluble 
in  an  excess  of  the  reagent. 

Manganese,  Mn 

Metallic  manganese  dissolves  easily  in  acids  with  the  for- 
mation of  the  manganous  ion,  Mn.  Manganese  is  also  found 
in  the  manganate  ion,  MnO"4,  which  is  green,  and  the  per- 
manganate ion,  Mn0'4,  which  is  purple. 


4O  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Use  for  the  reactions  of  the  manganous  ion  a  2-per-cent. 
solution  of  MnSO4. 

42.  The  sulphid  ion  (from  ammonium  sulphid)  precipitates 
the  manganous  ion  as  the  sulphid,  MnS,  a  flesh-colored  com- 
pound which  upon  standing  exposed  to  the  air  becomes  dark 
brown.     This  precipitate  is  insoluble  in  excess  of  the  reagent, 
but  dissolves  readily  in  acids. 

43.  The  hydroxid  ion  produces  a  precipitate  of  manganous 
hydroxid,  Mn(OH)2,  which  is  at  first  nearly  white,  but  when 
shaken  in  the  tube  with  air,  or  allowed  to  stand  exposed  to  it, 
becomes  quickly  dark  brown  from  the  absorption  of  oxygen 
and  the  formation  of  manganic  hydroxid,  Mn(OH)3.     The 
precipitate  is  insoluble  in  excess  of  sodium  hydroxid  but  dis- 
solves in  acids. 

If   ammonium  hydroxid  is  used  the  precipitation  is  pre- 
vented by  the  presence  of  ammonium  chlorid. 

44.  If  as  much  Pb3O4  (red  lead)  as  can  be  held  on  the  point 
of  a  pen-knife  is  added  to  a  solution  of  the  manganous  ion 
and  if  the  mixture  is  then  strongly  acidified  with  nitric  acid, 
boiled  and  allowed  to  settle,  the  liquid  above  has  a  red  color 
due  to  the  formation  of  the  permanganate  ion.     The  presence 
of  hydrochloric  acid  or  chlorin  compounds  interferes  with,  or 
prevents,  the  reaction. 


+  2PbSO4+i2H20,  or 


+  2PbSO4+i2H2O. 

45.  Solid  compounds  of  manganese  when  mixed  with 
several  times  their  weight  of  dry  sodium  carbonate  and  potas- 
sium nitrate  and  fused  on  the  platinum  foil  become  converted 
to  green  sodium  manganate,  NaaMnO*,  the  color  of  which  is 
best  seen  while  it  is  melted. 


+H20+CO2. 


METALS  (CATIONS)  41 

If  too  much  of  the  manganese  compound  is  used  the  mass 
may  be  black.  In  the  cold  it  dissolves  in  water  to  a  green 
solution.  This  is  decomposed  by  boiling  the  liquid,  the  manga- 
nese being  precipitated  as  an  oxid.  The  green  solution 
gradually  changed  to  a  pink  through  exposure  to  the  air,  the 
manganate  ion  being  converted  to  a  permanganate  ion,  MnO/. 

Iron,  Fe 

The  metal  is  soluble  in  hydrochloric,  sulphuric,  and  nitric 
acids.  There  are  two  series  of  compounds,  the  ferrous  and 
the  ferric.  The  former  are,  for  the  most  part,  colorless  or 
greenish  and  unite  with  oxygen  when  brought  in  contact  with 
oxidizing  agents,  or  even  when  exposed  to  the  air,  especially 
in  the  presence  of  water.  They  are  thus  converted  into  ferric 
compounds  which  are  usually  yellowish  to  reddish-brown. 

Ferrous  Ion  (Fe) 

For  the  reactions  use  a   2-per-cent.   solution  of  FeSC>4. 

46.  The  sulphid  ion  from  ammonium  sulphid  precipitates 
the  ferrous  ion  completely  from  neutral  or  alkaline  solutions 
as  black  ferrous  sulphid,  FeS.     This  dissolves  in  dilute  acids 
and  oxidizes  in  the  air  to  brown  ferric  hydroxid. 

47.  The  hydroxid  ion  produces  ferrous  hydroxid  Fe(OH)2. 
If  the  ferrous  solution  is  pure  and  oxygen  is  excluded  from 
that  and  the  reagent  the  color  at  first  is  nearly  white.     With- 
out these  special  precautions,  however,  it  is  a  grayish-green 
which  soon  changes  by  oxidation  to  a  darker  green  and  finally 
is   converted   to   reddish-brown   ferric   hydroxid,   Fe(OH)3. 
This  precipitation  is  prevented  by  the  presence  of  many  non- 
volatile organic  acids  or  their  salts  and  by  ammonium  salts; 
also  by  other  organic  compounds  like  sugar  and  glycerin. 

48.  The   ferrocyanid   ion,    Fe(CN)6///r,   (potassium  ferro- 
cyanid),  with  the  ferrous  ion  which  is  free  from  the  ferric 
ion  gives  a  bluish-white  precipitate  of  potassium  ferrous  fer- 
rocyanid, K2FeFe(CN)6.     This  becomes  deep  blue  by  the 


42  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

action  of  the  oxygen  of  the  air  or  by  oxidizing  agents.  Hy- 
droxids  of  the  alkalies  destroy  the  blue  color. 

49.  The  ferricyanid  ion,  Fe(CN)6"',  with  a  ferrous  solution 
produces  a  deep  blue  precipitate  of  ferrous  ferricyanid,  Fe3- 
Fe2(CN)i2.     Hydroxids    of    the    alkalies    decompose    this. 
With  very  dilute  solutions  of  iron  salts  the  last  two  reagents 
produce  only  a  greenish  color. 

50.  The  sulphocyanate  ion,  SCN',  imparts  no  color  to  fer- 
rous solutions  if  they  do  not  contain  ferric  compounds. 

5 1 .  The  carbonate  ion  precipitates  the  ferrous  ion  as  whi  te  fer- 
rous carbonate,  FeCO3,  which,  when  moist,  rapidly  oxidizes  in 
air,  giving  up  its  CO2,  ferric  hydroxid,  Fe(OH)3,  being  formed. 

52.  Tannic  or  gallic  acids  produce  no  color  with  ferrous 
solutions  if  they  are  free  from  ferric  salts. 

The  Ferric  Ion  (Fe) 

For   the   reactions   use   a    2-per-cent.  solution  of  FeCl3. 

53.  The  sulphid  ion  in  alkaline  solution,  e.g.,  from  am- 
monium sulphid,  precipitates  the  ferric  ion  as  black  ferric 
sulphid,  Fe2S3.     This,  for  the  most  part,  decomposes  upon 
acidifying,  with  reduction  of  the  iron  to  the  ferrous  condition 
and  precipitation  of  sulphur. 

Fe2S3+4H'Cl=2FeCl2+2H2S+S,  or 
Fe2S3+4lT  =2Fe+2H2S+S. 

54.  The  hydroxid  ion  precipitates  the  ferric  ion  as  ferric 
hydroxid,  Fe(OH)3,  a  reddish-brown,  gelatinous  substance. 
It  cannot  be  thus  formed  in  solutions  containing  a  large 
amount  of  non-volatile  organic  acids,  sugar  or  glycerin.     It 
dissolves  readily  in  acids,  but  is  insoluble  in  excess  of  the 
precipitant  and  in  solutions  of  ammonium  salts. 

Inasmuch  as  in  an  ionic  change  the  number  of  positive  and 
negative  charges  lost  must  be  eaual,  when  the  ferric  ion  loses 
six  of  the  positive  the  sulphur  ion  must  lose  the  same  number  of 
negative  charges.  Since  but  four  would  be  lost  from  the  two 


METALS  (CATIONS)  43 

sulphur  ions  which  unite  to  form  the  insoluble  ferrous  sulphid 
the  loss  of  two  more  negative  charges  causes  a  change  of  a 
sulphur  ion  to  the  elementary  form,  that  is,  the  sulphur  pre- 
cipitates. At  the  same  time  the  ion  has  changed  from  the 
ferric  to  the  ferrous  form,  that  is,  reduction  has  occurred. 

55.  The  ferrocyanid  ion  with  ferric  solutions  gives  a  dark 
blue  precipitate  of  ferric  ferrocyanid,  Fe4Fe3(CN)i8 — ''Prus- 
sian blue."     This  is  insoluble  in  hydrochloric  acid,  but  is  de- 
composed by  the  alkaline  hydroxids  with  the  separation  of 
ferric  hydroxid,  similarly  to  the  decomposition  of  ferrous 
ferrocyanid. 

56.  The  ferricyanid  ion  changes  the  color  of  ferric  solutions 
to  a  brown,  but  does  not  form  a  precipitate. 

57.  The  sulphocyanate  ion  gives  with  the  acidified  ferric 
ion  a  deep  blood-red. color  without  forming  a  precipitate. 
The  color  is  due  to  the  formation  of  non-ionized  ferric  sulpho- 
cyanate,  Fe(SCN)s.     An   excess  of  the  sulphocyanate  ion 
intensifies  the  color  by  decreasing  the  ionization  of  the  com- 
pound (p.  20).     Hydrochloric  acid  does  not  decolorize  it. 
The  reaction  may  be  used  as  an  extremely  sensitive  test  for 
iron.     When  soluble  phosphates,  borates  or  some  organic 
acids  are  present  the  reaction  may  fail  until  the  solution  is 
strongly  acidified  with  hydrochloric  acid.     The  color  is  de- 
stroyed by  the  addition  of  mercuric  chlorid. 

58.  The  carbonate  ion  precipitates  from  ferric  solutions 
reddish-brown,   gelatinous  ferric   hydroxid,   Fe(OH)3,    C02 
escaping. 

59.  Barium  carbonate  precipitates  ferric  compounds  com- 
pletely in  the  cold  as  basic  salts. 

60.  Tannic  or  gallic  acid  forms  a  blue-black  precipitate 
with  ferric  solutions. 

Nickel,  Ni 

Nickel  dissolves  slowly  in  hydrochloric  and  in  dilute  sul- 
phuric acids,  more  readily  in  nitric  acid.     Nickelous  salts  of 


44  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

these  acids  are  then  formed.     Most  of  the  solutions  of  nickel 
are  green,  the  color  of  the  nickel  ion. 

A  2-per-cent.  solution  of  NiSO4  may  be  used  for  the 
reactions. 

61.  The  sulphid  ion  from  ammonium  sulphid  in  neutral 
or  alkaline  solutions  precipitates  black  nickel  sulphid,  NiS, 
slightly  soluble  in  excess  of  the  reagent,  in  the  presence  of  am- 
monium hydroxid,  to  a  brownish  solution.     It  does  not  dis- 
solve in  cold  dilute  hydrochloric  acid,  but  dissolves  in  aqua 
regia. 

62.  The  hydroxid  ion  gives  at  first  a  greenish  precipitate 
of  nickelous  hydroxid,  Ni(OH)2,  which  dissolves  to  a  blue 
liquid  in  a  slight  excess  of  ammonium  hydroxid,.  through  the 
formation  of  a  complex  ion  of  nickel  and  ammonia,  but  does 
not  in  excess  of  sodium  or  potassium  hydroxid. 

63.  With  sugar  and  some  ions  of  organic  acids,  like  tartaric, 
the  nickel  ion  enters  into  combinations  with   which    the 
hydroxid  ion  or  the  sulphid  ion  forms  no  precipitate.     Thus, 
if  a  solution  of  sodium  potassium  tartrate  (Rochelle  salt)  is 
added  to  Ni  and  the  mixture  is  then  made  strongly  alkaline 
with  NaOH  addition  of  the  sulphid  ion  does  not  precipitate 
NiS,  but  gives  a  brown  solution. 

64.  Potassium  nitrite  does  not  precipitate  nickel  from  its 

solution  in  acetic  acid. 

Cobalt,  Co 

Cobalt,  like  nickel,  dissolves  in  acids,  with  the  formation 
of  cobaltous  salts.  These  are  usually  pink  to  red  when  they 
contain  water  of  crystallization  and  blue  when  anhydrous. 
The  dilute  solutions  are  red  also,  the  color  of  the  cobaltous  ion. 

Use  a  2-per-cent.  solution  of  Co(N03)2  in  the  reactions. 

65.  The  sulphid  ion,  from  ammonium  sulphid,  precipitates 
the  cobalt  ion  from  neutral  or  alkaline  solutions  as  the  black 
sulphid,  CoS.     It  is  insoluble  in  excess  and  in  cold  hydrochlo- 
ric acid  but  dissolves,  like  the  nickel  sulphid,  in  aqua  regia. 


METALS  (CATIONS)  45 

Tartrates  do  not  prevent  the  precipitation  of  CoS  from  alka- 
line solutions  as  they  do  of  NiS. 

66.  The  hydroxid  ion  gives  a  precipitate  of  a  blue  basic 
salt.     It  dissolves  in  excess  of  the  ammonia,  with  the  forma- 
tion of  a  complex  cobalt-ammonium  ion,  to  a  pink  solution  if 
the  liquid  is  freed  from  oxygen  and  this,  on  standing  in  the 
air,  becomes  brown  from  oxidation.     Organic  compounds 
may  prevent  the  precipitation. 

67.  Upon  boiling  with  an  excess  of  sodium  hydroxid  or  po- 
tassium hydroxid  the  blue  basic  salt  is  converted  into  pink 
insoluble  cobaltous  hydroxid,  Co(OH)2. 

68.  A  solution  of  cobalt  when  made  strongly  acid  with 
acetic  acid,  after  the  addition  of  a  considerable  amount  of 
potassium  nitrite  and  standing  for  several  hours  in  a  warm 
place,  will  give  a  yellow,  crystalline  precipitate  of  potassium 
cobaltic  nitrite,  K3Co(NO2)3. 

The  cobaltous  ion,  Co,  is  converted  into  the  cobaltic,  Co, 
by  oxidizing  agents,  e.g.,  nitrous  acid  which  is  set  free  through 
the  decomposition  of  the  nitrite  by  the  acetic  acid.  The 
complex  ion  Q^NC^V  is  thus  formed,  the  potassium  salt  of 
which  is  insoluble. 

Directions  for  the  Separation  of  Cations  of  Group  III  in 

the  Absence  of  Phosphoric,  Oxalic,  Boric,  Citric, 

and  Tartaric  Acids 

The  acidity  of  the  solution  containing  the  group  should  be 
neutralized  by  adding  ammonium  hydroxid  slowly,  until  a 
permanent  precipitate  begins  to  form.  Then  add  about  one- 
tenth  its  volume  of  ammonium  chlorid  and  last  ammonium 
sulphid  until  complete  precipitation  has  been  produced,  but 
avoiding  a  large  excess.  Heat  the  mixture  nearly  to  boiling 
and  filter.  If  the  original  mixture  contained  groups  IV  and  V 
they  will  be  found  in  the  filtrate.  Otherwise  it  can  be  dis- 
carded. The  precipitate  contains  the  sulphids  of  nickel, 


46  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

cobalt,  iron,  manganese,  and  zinc,  with  the  hydroxids  of 
chromium  and  aluminum.  Wash  it  with  hot  water. 

Rinse  the  precipitate  from  the  filter  with  about  15  c.c.  of 
dilute  hydrochloric  acid  and  let  it  stand  five  minutes  in  the 
acid,  stirring  occasionally.  All  the  members  of  the  group 
with  the  exception  of  nickel  and  cobalt  dissolve.  The  nickel 
aod  cobalt  remain  as  black  sulphids.  If  the  insoluble  residue 
is  light  colored  or  gray  it  may  be  only  sulphur.  Filter  and 
wash. 

Test  the  residue  with  the  borax  bead  in  the  oxidizing  flame. 
Cobalt  gives  a  blue  bead;  nickel  a  purple,  white  hot,  soon 
becoming  brown.  II  the  residue  is  so  small  that  it  cannot  be 
removed  from  the  paper,  the  part  of  the  filter  which  contains 
it  can  be  burned  on  the  bead,  when,  after  the  complete  oxida- 
tion of  the  carbon,  the  metal  will  be  left  to  dissolve  in  the 
borax. 

When  only  one  of  the  two  is  present  the  colors  will  be  clear 
but  if  both  are  contained  in  the  bead,  it  may  be  necessary  to 
separate  them  before  they  can  be  identified. 

One  method  to  accomplish  this  is  to  dissolve  the  mixed  sul- 
phids in  2  c.c.  of  warrrfaqua  regia,  filter  and  evaporate  nearly 
to  dryness  under  a  hood,  to  remove  the  excess  of  acid.  Dilute 
with  water  to  5  c.c.,  add  potassium  nitrite  and  acidify  strongly 
with  acetic  acid;  warm  gently  and  let  it  stand  several  hours. 
The  cobalt  gives  a  yellow  precipitate.  Filter  and  test  the  fil- 
trate for  nickel  by  sodium  hydroxid  or  ammonium  sulphid. 
The  former  gives  a  green  precipitate;  the  latter,  a  black  one. 

As  an  alternate  method  of  separation  of  cobalt  and  nickel, 
after  dissolving  the  sulphids  in  i  c.c.  of  hydrochloric  acid  with 
0.5  c.c.  of  nitric  acid  add  a  solution  of  Rochelle  salt;  make 
alkaline  with  ammonium  hydroxid  and  add  an  excess  of  am- 
monium sulphid;  black  cobalt  sulphid  is  precipitated;  nickel 
remains  dissolved  as  a  brown  compound,  seen  after  filtration. 

Boil,  for  a  short  time,  the  hydrochloric  acid  solution  of  iron, 


METALS  (CATIONS)  47 

manganese,  chromium,  aluminum,  and  zinc  until  it  is  freed 
from  hydrogen  sulphid  as  shown  by  the  absence  of  odor  or  by 
its  failure  to  discolor  a  paper  dipped  in  lead  acetate  and  held 
in  the  steam.  Add  to  the  hot  liquid  bromin  water  until  it  is 
colored  yellow  and  bring  to  a  boil.  Make  it  alkaline  with  so- 
dium hydroxid,  then  add  enough  of  the  sodium  hydroxid  to 
make  a  large  excess.1  Boil  one  minute  and  filter.  The  iron, 
manganese,  and  chromium  are  precipitated  as  hydroxids,  and 
the  zinc  and  aluminum  are  dissolved  in  the  excess  of  the  alkali. 
Dissolve  a  small  portion  of  the  precipitate  in  hydrochloric 
acid  and  test  it  for  iron  with  potassium  ferrocyanid.  A  blue 
color  is  produced.  Place  another  portion  of  the  precipitate 
on  the  platinum  foil,  add  five  times  as  much  of  a  mixture  of 
dry  sodium  carbonate  and  potassium  nitrate  and  heat  until  it 
is  thoroughly  fused.  If  manganese  is  present  the  mass  is  a 
deep  green.  Chromium,  in  the  absence  of  manganese,  gives 
a  yellow  mass.  In  either  case,  to  confirm  the  presence  of 
chromium  place  the  foil  in  a  test-tube,  cover  with  water  and 
heat  to  boiling.  Filter  off  the  brownish  residue,  if  one  is  pres- 
ent, and  test  the  solution  for  chromium,  after  acidifying  with 
acetic  acid,  by  adding  a  few  drops  of  lead  acetate.  A  yellow^ 
precipitate  of  lead  chromate  is  produced. 

The  solution  of  aluminum  and  zinc  in  excess  of  sodium 
hydroxid  should  be  slightly  acidified  with  hydrochloric  acid, 
then  an  excess  of  ammonium  hydroxid  should  be  added  and 
the  liquid  boiled  for  several  minutes.  Aluminum  hydroxid, 
a  very  flocculent  white  precipitate,  is  produced,  which  floats 
a  long  time  in  the  liquid.  Filter  and,  after  dividing  the  fil- 
trate into  two  parts,  test  one  for  zinc  by  a  few  drops  of  ammo- 
nium sulphid  and  warming.  A  flocculent  white  precipitate 
appears.  From  the  second  part,  after  it  has  been  acidified 
with  hydrochloric  acid,  potassium  ferrocyanid  precipitates 
zinc  ferrocyanid,  a  very  fine,  white  solid. 

1  When  sufficient  alkali  has  been  added  the  liquid  has  a  soapy  feeling  if  a 
drop  js  rubbed  between  the  fingers. 


INTRODUCTION   TO    CHEMICAL   ANALYSIS 

TABLE  II 

OUTLINE  OF  SEPARATION  or  THE  CATIONS  OF  GROUP  III  IN  ABSENCE  OF 
/  PHOSPHORIC,  BORIC,  ETC.,  ACIDS 

To  the  solution  add  NH4OH  until,  after  shaking,  there  is  a  slight  perma- 
nent precipitate.,  Then  add  NH4C1  and  a  slight  excess  of  (NH4)2S.  Warm, 
filter,  and  washy  The  precipitate  contains  NiS,  CoS,  FeS  (or~Fe2S3),  MnS, 
ZnS,  Cr(OH)3,  and  A1(OH)3  if  they  were  present  in  the  original  solution. 
Let  it  stand  five  minutes  in  cold  dilute  HC1,  stirring  frequently.  All  are 
dissolved  except  the  sulphids  of  Ni  and  Co.  Filter  and  wash. 


the  insoluble  residue 
Is  black  apply  the  borax 
bead  test. 

A  blue  bead  =  Co. 

A  brown  bead  =  Ni. 

If  a  separation  of  the 
metals  is  desired  dissolve  ! 
the  black  residue  in  a 
little  aqua  regia,  evapo-  * 
rate  nearly  to  dryness, 
dilute  with  5  c.c.  of 
water,  add  KNO2, 
acidify  with  acetic  acid 
and  let  it  stand  several 
hours,  after  warming.  A 
yellow  precipitate  of  po- 
tassium cobalt  nitrite  is 
formed.  Filter  and  to 
the  filtrate  add  NaOH 
and  warm.  Ni(OH)2  a 
light  green  precipitate 
appears.  With  the  same 
filtrate  (NH4)2S  gives  a 
black  precipitate,  NiS. 

Or,  add  excess  of 
NaOH  to  aqua  regia  so- 
lution, then  a  tartrate 
and  (NH4)2S.  Ni  gives  a 
brown  solution;  ACo,  a 
black  precipitate. 


filtrate  contains  Fe,  Mn,  Zn,  Cr,  and  Al 


Expel  the  hydrogen  sulphid  by  boiling  and  add  a 
slight  excess  of  bromin  water. 

Then  heat  to  boiling,  make  the  liquid  very 
strongly  alkaline  with  NaOH,  and  boil  one  minute. 
Filter. 


he  precipitate  contains' 
Fe(OH)3,  Cr(OH)3,  and 
Mn(OH)2. 

Dissolve  a  small  portion 
of  the  precipitate  in  HC1 
and  test  for  Fe  with  K4Fe- 
(CN)6.  A  dark  blue  color 
should  appear. 

Dry  another  portion  of 
the  precipitate  on  plati- 
num foil,  and  thoroughly 
fuse  with  Na2CO3  and 
KN03. 

Dark  green  K2Mn04  is 
formed. 

Boil  with  a  lij&le  water, 
filter,  acidify  wafeacetic 
acid  and  add  lea<T acetate. 
Yellow  PbCrO?  is  pre- 
cipitated. 


^The  filtrate  contains 
Al  and  Zn. 

Slightly  acidify  it 
with  HC1,  make  it 
alkaline  with  NH4OH 
and  boil  a  minute.  A 
light  white  precipitate 
is  A1(OH)3. 

Filter  and  test  the 
filtrate  for  Zn. 

1.  (NH4)2S  gives  a 
white    flocculent    pre- 
cipitate, ZnS. 

2.  K4Fe(CN>6  gives 
a  fine  white  precipitate 
of    Zn2Fe(CN)6    after 
acidifying  with  HC1. 


Explanation  of  the   Operations  Used  in  the   Separation 
of  the  Cations  of  Group  III 

The  object  of  the  addition  of  ammonium  hydroxid  is  to 
neutralize  any  acid  present,  since  this  would  otherwise  decom- 
pose the  ammonium  sulphid  and  prevent  the  precipitation  of 
the  cations.  Ammonium  chlorid  favors  the  complete  precipi- 


METALS  (CATIONS)  49 

tation  of  many  of  these,  converting  the  products  from  a  gelati- 
nous to  a  more  or  less  flaky  consistency,  which  hAstens  the 
subsequent  filtration.  Warming  likewise  favors  tne  precipi- 
tation and  rapid  filtration  of  the  group.  The  nfrecipi tation 
should  be  accomplished  chiefly  by  the  ammonium  sulphid,  a 
large  excess,  of  which,  together  with  ammonium  hydroxid 
may  cause  a  partial  solution  of  the  nickel  sulphid.  This  gives 
the  filtrate  a  brown  color. 

By  the  action  of  the  hydrochloric  acid  the  compounds  of  all 
the  metals,  with  the  exception  of  nickel  and  cobalt,  are  con- 
verted to  chlorids  which  are  soluble.  The  sulphids  of  these 
two  metals  are  not  soluble  in  cold  dilute  acid,  although  they 
are  somewhat  so  if  it  is  heated  or  in  the  concentrated  acid. 
Sulphur  is  set  free  from  the  ammonium  sulphid  by  the  action 
of  ferric  salts  (53). 

When  the  sulphids  of  the  metals  are  dissolved  in  the  acid 
hydrogen  sulphid  is  formed,  and  remains  dissolved  in  the  cold 
liquid.  If  it  is  not  removed  it  would,  as  soon  as  the  liquid 
becomes  alkaline  in  the  next  operation,  change  the  iron,  man- 
ganese, and  zinc  compounds  again  into  the  sulphids  instead 
of  the  hydroxids.  The  zinc  would  remain  with  them, 'in- 
stead of  being  separated  by  the  sodium  hydroxid. 

Bromin  water  and  other  oxidizing  agents  change  ferrous 
compounds  to  ferric  in  the  presence  of  free  acid.  After  this 
has  been  added  it  is  necessary  to  use  a  large  excess  of  the 
sodium  hydroxid,  since  by  a  small  amount  all  five  of  the 
cations  in  the  solution  would  be  precipitated? 

The  fusion  of  the  hydroxids  of  iron,  manganese,  and  chro- 
mium does  not  change  the  appearance  of  the  iron  compound, 
but  converts  the  manganese  into  a  green  manganate  which  is 
changed  by  boiling  into  an  insoluble  oxid  (45).  At  the  same 
time  the  chromium  hydroxid  is  oxidized  to  a  soluble  yellow 
chromate  (37). 

If  the  aluminum  solution  is  not  made  acid  the  aluminum 

4 


50  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

will  not  be  precipitated  by  ammonium  hydroxid  even  if  large 
amounts  are  present  (33). 

Separation  of  the  Cations  of  Groups  HI  and  IV  in  the 
Presence  of  Phosphoric  Acid  or  its  Salts 

(Capital  and  small  letters  are  used  for  the  identification  of  subdivisions  of  the 
groups — the  former  for  solids,  the  latter  for  solutions) 

The  cations  of  group  III  and  IV  are  ordinarily  precipitated 
separately  from  those  of  group  V.  If,  however,  the  solution 
contains  certain  acids  or  their  salts  the  cations  of  both  groups 
are  thrown  down  together  by  the  reagents  used  for  the  pre- 
cipitation of  group  III.  If  it  has  been  ascertained  by  testing 
that  phosphates  are  present,  the  usual  separation  of  the  ca- 
tions of  groups  III  and  IV  should  be  modified  in  the  following 
manner: 

To  the  solution  of  the  two  groups  add  about  one-tenth  its 
volume  of  ammonium  chlorid,  make  it  faintly  alkaline  with 
ammonium  hydroxid  and  then  add  ammonium  sulphid  until 
no  more  precipitate  (A)  forms.  Warm  and  stir  the  mixture 
until  the  precipitate  begins  to  settle,  filter  and  wash  with 
water  containing  a  few  drops  of  (NH^S.  Preserve  the  fil- 
trate to  test  for  cations  of  group  IV  which  may  be  partially 
unprecipitated.  The  precipitate  (A)  may  contain  all  the 
metals  of  groups  III  and  IV,  together  with  some  of  V  and  sul- 
phur. Rinse  it  from  the  paper  with  about  15  c.c.  of  dilute 
hydrochloric  ajid  into  a  dish  or  beaker  and  stir  it  occasionally 
for  five  minutes.  Filter  and  wash.  There  will  remain  in- 
soluble (B)  the  sulphids  of  cobalt  and  nickel  with  sulphur. 
The  remaining  compounds  are  dissolved  and  passed  into  the 
filtrate  (a). 

Nickel  and  cobalt  sulphids,  if  present,  are  black;  sulphur  is 
grayish  to  yellow,  burning  with  characteristic  odor  and  blue 
flame  is  present  in  large  quantity.  To  test  for  the  nickel  or 
cobalt  dissolve  the  residue  (B)  in  a  borax  bead  with  the  aid  of 


METALS  (CATIONS)  51 

the  oxidizing  flame.  Cobalt  gives  a  deep  blue  color;  nickel, 
a  purple  while  hot,  quickly  passing  to  a  brown.  In  cases  of 
doubt  or  where  it  is  desirable  to  separate  the  metals,  dis- 
solve it  in  aqua  regia  and  use  potassium  nitrite  or  tartrate  as 
on  page  46. 

The  hydrochloric  acid  solution  (a)  is  to  be  used  for  the 
separation  of  the  remaining  cations.  Boil  it  in  a  beaker  or 
flask  until  the  hydrogen  sulphid  is  expelled,  as  shown  by  the 
absence  of  odor  and  by  the  failure  of  a  strip  of  filter-paper  to 
turn  brown  after  it  has  been  moistened  with  lead  acetate  solu- 
tion and  suspended  in  the  steam.  If  much  sulphur  sepa- 
rates during  this  boiling  filter  it  out. 

Mix  about  one-fourth  of  the  solution  (a)  with  dilute  sul- 
phuric acid.  The  sulphate  of  barium  will  be  immediately 
precipitated  (C),  that  of  strontium,  if  much  is  present,  or  upon 
standing,  and  of  calcium  if  this  is  present  in  concentrated 
solution,  although  the  calcium  in  great  part  remains  dissolved. 
After  it  has  settled  filter  and  wash,  preserving  the  filtrate  (b) . 
If  there  is  no  precipitation  by  sulphuric  acid  calcium  may 
nevertheless  be  in  solution.  To  detect  it  here  or  after  filtra- 
tion from  the  sulphates  of  barium  and  strontium  add  three 
volumes  of  alcohol  to  (b)  when  the  calcium  sulphate  will  be 
precipitated  if  present.  It  can  be  dissolved  in  boiling  water 
and  tested  with  a  few  drops  of  ammonium  oxalate  which 
should  give  a  fine  white  precipitate  of  calcium  oxalate.  To 
test  for  barium  and  strontium  if  sulphuric  acid  caused  a  pre- 
cipitation place  the  washed,  moist  precipitate  (C)  in  a  porce- 
lain dish,  add  about  10  c.c.  of  a  ic-per-cent.  solution  of 
sodium  carbonate,  and  boil  gently  for  five  minutes.  Stron- 
tium sulphate,  and  calcium  sulphate  if  it  is  present,  are  con- 
verted into  carbonates,  the  barium  sulphate  being  but  little 
affected.  Filter  and  wash.  Pour  over  the  precipitate  (D) 
on  the  filter  10  c.c.  of  hot  acetic  acid  which  will  dissolve  (c) 
the  strontium  carbonate,  but  comparatively  little  of  the 


52  INTRODUCTION"  TO    CHEMICAL   ANALYSIS 

barium  salt.  Wash  the  precipitate  (D) ,  discarding  the  wash- 
water.  The  barium  may  be  confirmed  by  the  appearance  of  a 
yellow-green  flame  obtained  from  the  insoluble  residue  (D) 
after  it  has  been  moistened  with  hydrochloric  acid.  To  the 
acetic  acid  solution  (c)  add  a  few  drops  of  potassium  dichro- 
mate  to  precipitate  the  barium  if  any  of  this  has  been  dis- 
solved. This  will  be  changed  to  yellow  barium  chromate. 
Filter  when  enough  of  the  reagent  has  been  added  to  give  the 
supernatant  liquid  a  yellow  color.  From  the  nitrate  (d)  pre- 
cipitate the  strontium  by  means  of  dilute  sulphuric  acid,  add- 
ing it  as  long  as  the  insoluble  compound  is  formed  and  letting 
the  mixture  stand  five  minutes  if  no  precipitate  appears  imme- 
diately. Filter  and,  after  washing,  confirm  the  strontium 
by  moistening  with  hydrochloric  acid  and  making  the  flame 
test  with  and  without  the  blue  glass  (19). 

To  about  a  cubic  centimeter  of  the  hydrochloric  acid  solu- 
tion (a)  add  enough  bromin  water  to  give  a  yellow  color,  boil 
and  add  a  few  drops  of  potassium  ferrocyanid.  A  blue  color 
is  indicative  of  iron. 

To  the  remainder  of  the  hydrochloric  acid  solution  (a),  if 
iron  has  been  found,  add  bromin  water  and  boil  as  before. 
If  iron  is  absent  omit  this  step  and  proceed  with  the  next 
operation,  which  is  to  remove  from  the  solution  (a)  the  phos- 
phoric acid  that  may  be  present.  This  is  effected  by  forming 
an  insoluble  compound  of  this  with  iron.  To  learn  whether 
enough  iron  is  present  in  the  solution  remove  two  or  three 
drops  and  to  these,  in  an  evaporating  dish,  add  ammonium 
hydroxid.  If  there  is  a  yellowish-brown  precipitate  the 
amount  of  iron  is  sufficient;  if  it  is  nearly  white  add  to  the 
solution  a  few  drops  of  ferric  chlorid,  and  test  a  little  of  this 
with  ammonium  hydroxid  as  before,  repeating  as  many  times 
as  necessary  to  produce  the  yellow  or  brown  precipitate. 
The  free  acid  should  now  be  nearly  but  not  quite  neutralized. 
This  must  be  done  by  dropping  in,  with  constant  stirring, 


METALS  (CATIONS)  53 

sodium  hydroxid  solution  until  blue  litmus-paper  is  only 
slowly  turned  red.  If  too  much  has  been  used  hydrochloric 
acid  must  be  added  until  the  proper  reaction  is  attained. 
Now  add  an  excess  of  barium  carbonate.  This  is  best  done 
by  shaking  the  solid  reagent  with  water  and  using  the  milky 
mixture.  Let  this  stand  in  the  cold  for  half  an  hour,  with  fre- 
quent stirring.  The  liquid  above  the  precipitate  should  be 
colorless  if  enough  has  been  used.  Then  filter  and  wash. 
The  precipitate  (E)  contains  the  hydroxids  of  aluminum, 
chromium,  and  iron,  also  the  ferric  phosphate  and  the  excess 
of  the  barium  carbonate.  The  filtrate  (e)  will  contain  the 
manganese  and  zinc,  also  the  barium,  strontium,  calcium, 
and  magnesium,  which  may  have  been  contained  in  the  origi- 
nal substance. 

Transfer  the  precipitate  (£)  from  the  filter  to  a  porcelain 
dish,  then  boil  it  two  minutes  with  about  10  c.c.  of  sodium 
hydroxid,  cool  and  filter.  This  dissolves  the  aluminum, 
which  passes  into  the  filtrate.  From  this  it  can  be  precipi- 
tated by  first  acidifying  with  hydrochloric  acid,  then  making 
faintly  alkaline  with  ammonium  hydroxid  and  boiling  for  a 
minute  in  a  beaker.  The  aluminum  hydroxid  thus  obtained 
is  a  white,  very  voluminous  substance,  which  floats  for  a  long 
time  in  the  liquid. 

To  detect  the  chromium  in  the  portion  of  the  barium  car- 
bonate precipitate  (E)  which  did  not  dissolve  in  sodium 
hydroxid,  dry  it  on  platinum  foil,  then  fuse  thoroughly  with 
a  mixture  of  dry  sodium  carbonate  and  potassium  nitrate. 
The  chromium  produces  a  yellow  mass  which  dissolves  to  a 
yellow  solution  in  water,  and  this,  after  acidifying  with  acetic 
acid,  forms  a  bright  yellow,  insoluble  lead  chromate  upon  the 
addition  of  lead  acetate. 

To  test  for  magnesium,  zinc,  and  manganese  proceed  as 
follows  with  the  filtrate  (e)  from  the  barium  carbonate  pre- 
cipitate. Acidify  with  dilute  sulphuric  acid,  heating  until 


54  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

the  effervescence  has  ceased.  This  precipitates  the  barium, 
most  of  the  strontium,  and  possibly  part  of  the  calcium,  and 
dissolves  the  manganese,  zinc,  and  magnesium.  Filter  and 
neglect  the  precipitate,  since  the  bases  contained  in  it  have 
already  been  determined.  To  the  nitrate  add  ammonium 
hydroxid  until  it  is  alkaline,  then  ammonium  sulphid.  Man- 
ganese and  zinc  are  precipitated  (F)  as  sulphids  if  they  were 
contained  in  the  mixture.  Magnesium,  with  calcium  and 
possibly  some  strontium,  remains  in  the  liquid  (/).  Filter 
and  wash.  Remove  the  precipitate  (F)  from  the  paper,  dis- 
solve it  in  a  slight  excess  of  hydrochloric  acid,  free  it  from 
hydrogen  sulphid  by  boiling  in  an  evaporating  dish,  and  make 
it  strongly  alkaline  with  sodium  hydroxid.  White  mangan- 
ous  hydroxid,  soon  turning  brown,  is  precipitated,  the  zinc 
remaining  in  solution  (g) .  Confirm  the  manganese  by  fusing 
it  on  the  platinum  foil  with  sodium  carbonate  and  potassium 
nitrate,  when  a  dark  green  mass  results.  The  zinc  may  be 
separated  from  the  solution  (g)  in  sodium  hydroxid  by  the  aid 
of  a  little  hydrogen  sulphid,  which  gives  a  white  flaky  precipi- 
tate of  zinc  sulphid. 

To  determine  whether  any  magnesium  is  contained  in  the 
solution  (/)  from  which  the  manganese  and  zinc  sulphids  have 
been  filtered,  add  ammonium  oxalate  which  may  give  a  white 
precipitate  of  calcium  or  strontium  oxalate.  Remove  this  by 
filtration  if  it  is  formed  and  to  the  filtrate  add  ammonium 
hydroxid  and  sodium  phosphate.  A  white  crystalline  pre- 
cipitate indicates  magnesium. 


METALS  (CATIONS) 


55 


TABLE  III 

OUTLINE   OF    SEPARATION  or  THE  CATIONS  OF  GROUPS  III  AND  IV 

In  the  presence  of  phosphoric  acid  or  its  salts.  To  the  solution  add 
NH4C1,  make  alkaline  with  NH4OH  and  precipitate  with  (NH4)2S.  Warm, 
filter,  and  wash.  The  precipitate  contains  NiS,  CoS,  FeS  (or  Fe2S3),  MnS, 
ZnS,  A1(OH)3,  and  Cr(OH)3,  with  Ba,  Sr,  Ca,  and  Mg  salts  of  phosphoric 
acid,  sometimes  with  S.  Treat  with  HC1.  If  a  residue  remains,  filter  and 
wash. 


If  the 

The  nitrate  contains  Fe,  Mn,  Cr,  Al,  Zn,  Ba,  Sr,  Ca,  and  Mg.     Expel  the 

residue  is 

H2S  by  boiling  and  divide  the  solution  into  three  portions.     To  the  first  add 

black    it 
nitLV  con." 

dilute  H2SO4.     Filter  and  wash. 

tain  CoS 

Precipitate 

Filtrate 

Filtrate  contains  Fe,  Mn,  Cr,  Al,  Zn,  Ca,  and 

and  NiS. 

contains 

contains 

Mg.     Test  a  portion  for  Fe  by  K4Fe(CNj6.     If 

Test     it 

BaSO4      and 

Ca.       Mix 

iron  is  present,  after  oxidizing  this  with  Br  water 

with  b  o- 

SrSO4       with 

a       second 

add  enough  FeCl3  to  unite  with  the  phosphoric 

rax  bead. 
A  blue 

possibly 
CaSO4.     Boil 

portion 
with  3  vol- 

acid, nearly  neutralize  the  acidity  and  add  BaCO3. 
Allow   to   stand   one-half   hour.     Then    filter  and 

bead  in- 

with Na2CO3, 

umes  of  al- 

wash. 

dicates 

filter     ciiicl 

cohol      A 

Co,     a 
brown 
bead  in- 
dicates 

•vr:        Tr\ 

wash.          In- 
soluble    resi- 
due   contains 
BaSO4 

white    pre- 
cipitate  of 
CaSO4. 
Dissolve  in 

Precipitate  contains 
Fe(OH)3,       Cr(OH)3, 
and     A1(OH)3.      Boil 
with  NaOH  and  filter. 

Filtrate  contains  Mn,  Zn, 
Ca(Sr?),  and  Mg.     Acidify 
with     dilute     H2SO4    and 
filter  out  the  Ba  and  Sr  if 

INI.        1O 

separate 
or      dis- 
tinguish 
these  dis- 

(BaC03?), 
SrCO3 
(CaC03?). 
Treat       with 
acetic      acid. 

hot     water 
and  confirm 
by  (NH4)2- 
C204. 

there  is  a  precipitate.     In 
the  filtrate  precipitate  the 
Mn  and  Zn  with  NH4OH 
and  (NH4)2S.     Filter. 

Residue 
contains 
Cr(OH),. 

Filtrate 
contains 
Al.    Acid- 

solve    in 
aqua 
regia  and 
use 

Insoluble 
BaSO4        re- 
mains.    Con- 

•£*••**•                              •«-rrt4-'U 

Fuse  with 
Na2CO3 
and 
KNO3. 

ify      with 
HC1,    add 
NH4OH 
and    boil. 

Precipitate 
contains  Mns 
and  ZnS. 

Filtrate  con- 
tains       (Ba, 
Sr?)   Ca  and 

KN02  or 

nrm         witn 

{\n-rmn                     4-f*c*i- 

Dissolve 

A       light 

Dissolve    in 

Mg.       With 

Rochelle 
salt  as  on 
page  46. 

name        test. 
From  solution 
precipitate  Ba 
with  K2Cr2O7 
^filter  and  add 
H2SO4;  white 
precipitate  of 
SrSO4.    Filter, 

in    water, 
add  acetic 
acid     and 
lead    ace- 
tate.       A 
yellow  pre- 
cipitate of 
PbCrO4. 

white  pre- 
cipitate of 
A1(OH)3. 

HC1.    Expel 
H2S  by  boil- 
ing and  add 
large    excess 
of    NaOH. 
Mn(OH)2  is 
precipitated. 
Confirm    by 

(NH4)2C2O4 
precipitate 
all  except  Mg 
and    remove 
by  filtration. 
To    the    fil- 
trate       add 
Na2HPO4. 

wash,     and 
confirm     by 
crimson  flame. 

fusion    with 
Na2CO3  and 
KNO3  on  Pt 

Mg    gives    a 
white      crys- 
talline     pre- 

foil.         Mn 

cipitate       of 

gives  a  green 

NH4MgPO4. 

mass.     From 

filtrate    pre- 

cipitate 
white,   flaky 

ZnS  by  H2S. 

56  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

Explanation  of  the  Operations  Used  in  the  Separation  of 

the  Cations  of  Groups  III  and  IV  in  the  Presence 

of  Phosphoric  Acid  or  its  Salts 

The  reason  that  it  is  necessary  to  modify  the  usual  methods 
when  this  acid  is  contained  in  the  solution  is  that  the  phos- 
phates of  barium,  strontium,  calcium,  and  magnesium  are 
soluble  in  acids,"  bul  not  in  neutral  or  alkaline  liquids.  They 
therefore  will  be  precipitated  with  the  third  group  as  soon 
as  ammonium  hydroxid  has  been  added.  Many  of  the  ex- 
planations given  regarding  the  analysis  of  groups  III  and  IV 
will  also  be  applicable  here. 

Nickel  and  cobalt  can  be  easily  separated  from  the  rest 
of  the  precipitated  metals  because  of  the  insolubility  of  their 
sulphids  in  dilute  acid.  From  the  solution  thus  obtained  all 
the  barium  and  most  of  the  strontium  are  separated  by  con- 
version into  sulphates  by  means  of  sulphuric  acid  which  dis- 
solves the  other  compounds.  This  does  not  make  a  clean 
separation  between  the  alkali  earth  metals,  though,  since  a 
part  of  the  strontium  may  be  left  in  solution  or  a  part  of  the 
calcium  may  be  changed  to  the  insoluble  sulphate  if  a  large 
amount  is  present.  Nevertheless  it  serves  to  distinguish 
them  from  each  other.  To  give  greater  certainty  of  this  the 
flame  test  should  not  be  neglected.  In  the  same  way  when 
the  sulphates  of  barium  and  strontium  are  boiled  with  sodium 
carbonate  some  of  the  barium  sulphate  may  become  con- 
verted to  the  carbonate  as  the  strontium  sulphate  is  and  thus 
pass  with  the  strontium  into  solution. 

The  separation  of  the  metals  contained  in  the  sulphuric 
acid  nitrate  is  the  same  as  in  the  separation  where  the  metals 
of  the  alkaline  earths  are  not  present,  except  that  alcohol  is 
used  to  render  more  insoluble  the  calcium  sulphate,  and  after 
the  metals  of  group  III  have  been  identified  and  the  calcium 


METALS  (CATIONS)  57 

removed  the  magnesium  is  precipitated  by  sodium  phosphate 
in  the  usual  manner. 

Practical  Exercises  on  Group  III 

1.  In  a  large  test-tube  mix  together  3-4  c.c.  of  the  solutions 
of  each  of  the  cations  of  the  group  and  separate  them  accord- 
ing to  Table  II. 

2.  In  the  same  manner  make  analyses  of  the  unknown  mix- 
tures which  will  be  furnished  by  the  instructors. 

3.  To  solutions  of  these  metals  add  those  of  group  IV  and 
sodium  phosphate  and  make  the  analysis  by  the  directions  in 
Table  III. 

4.  Examine  the  unknowns  furnished  by  the  instructors  to 
determine  whether  they  contain  phosphates.     If  this  is  the 
case  analyze  the  mixture  by  Table  III. 

Questions  for  Further  Study  on  Group  ffl 

In  the  fusion  of  chromium  and  manganese  compounds  with 
sodium  carbonate  and  potassium  nitrate  what  action  has  each 
of  the  latter  reagents?  What  is  the  action  of  the  PbaC^  and 
nitric  acid  in  the  production  of  the  red  color  from  manganous 
compounds?  Would  ferrous  or  ferric  salts  be  formed  when 
iron  is  dissolved  in  hydrochloric  acid?  in  dilute  sulphuric 
acid?  in  concentrated  nitric  acid?  Which  will  the  more 
easily  dissolve  iron,  concentrated  or  dilute  sulphuric  acid? 
What  are  some  of  the  pharmaceutical  preparations  of  iron 
from  which  ammonium  hydroxid  would  fail  to  produce  a 
precipitate?  What  is  the  valence  of  the  iron  in  ferrocyanid 
and  f erricyanid  ?  For  what  is  ferrous  carbonate  used  in  medi- 
cine and  what  means  is  then  adopted  to  prevent  its  oxidation  ? 
Is  the  rapidity  of  oxidation  of  ferrous  compounds  modified 
by  the  presence  or  absence  of  moisture?  What  pharma- 
ceutical preparations  would  be  incompatible  with  iron  com- 
pounds because  of  their  containing  tannic  or  gallic  acid? 


,58  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

What  compounds  of  the  alkali  metals  are  chemically  in- 
compatible with  salts  of  this  group  because  of  the  resulting 
insoluble  compounds?  What  can  be  used  to  prevent  this 
precipitation?  What  chemical  reagents  can  be  employed  to 
convert  ferrous  into  ferric  compounds?  What  application  is 
made  of  the  mutual  decomposition  of  alum  and  an  alkaline 
carbonate. 

GROUP  II 

Arsenic,  antimony,  tin,  gold,  platinum,  mercury  (ic  com- 
pounds), lead,  bismuth,  copper,  cadmium.  The  cations  of 
this  group  are  precipitated  from  their  acid  solutions  by  hy- 
drogen sulphid.  For  convenience  they  may  be  divided  into 
two  divisions :  (A)  Those  of  which  the  sulphids  are  insoluble 
in  yellow  ammonium  sulphid — mercuric  compounds,  lead,  bis- 
muth, copper,  and  cadmium,  and  (B)  those  of  which  the  sul- 
phids are  soluble  in  yellow  ammonium  sulphid — arsenic, 
antimony,  tin,  gold,  and  platinum. 

Division  A 

The  sulphids  of  this  division  do  not  form  soluble  sulpho- 
salts  with  the  alkaline  sulphids  and  are  consequently  in- 
soluble in  these  reagents.  Their  oxids,  hydroxids,  carbonates, 
phosphates,  arsenates,  arsenites,  and  iodids  as  well  as  sul- 
phids are  insoluble  in  water. 

Copper,  Cu 

The  metal Js  practically  insoluble  in  dilute  sulphuric  or 
hydrochloric  acids,  but  dissolves  readily  in  nitric  acid.  Its 
soluble  compounds  are  poisonous. 

For  the  reactions  a  2-per-cent.  solution  of  CuS04  may  be 
used. 

69.  The  sulphid  ion  gives  a  brownish,  nearly  black,  pre- 
cipitate of  copper  sulphid,  CuS.  It  dissolves  in  hot  nitric 
acid,  also  in  solutions  of  potassium  cyanid. 


METALS  (CATIONS)  59 

70.  The  hydroxid  ion,  with  the  copper  ion,  forms  the  light 
blue,  insoluble,  cupric  hydroxid  Cu(OH)2,  which  is  insoluble 
in  excess  of  sodium  hydroxid  or  potassium  hydroxid.     It  is 
decomposed  by  boiling  into  black  cupric  oxid,  CuO,  and 
water.     The    presence    of    some    organic    compounds    like 
glycerin,  sugar,  and  salts  of  tartaric  and  some  other  organic 
acids  prevents  the  precipitation  of  the  copper  ion  by  the 
hydroxid  ion,  but  instead  it  forms  a  deep  blue  solution.     If 
one  of  the  reducing  sugars,  like  glucose,  is  present  this,  by 
boiling,  reduces  the  cupric  to  a  cuprous  compound  and  yellow 
or  reddish  cuprous  oxid  is  precipitated. 

71.  Ammonium  hydroxid  in  small  amounts  forms  with  the 
copper  ion  a  blue  precipitate,  but  this  dissolves  easily  in  an 
excess  of  the  reagent  to  a  deep  blue  solution  which  contains 
the  complex  ion,   (NHs^Cu.     The  addition  of  potassium 
cyanid  decolorizes  the  liquid. 

72.  The  carbonate  ion  precipitates  basic  cupric  carbonate 
which  dissolves  in  excess  of  ammonium  carbonate. 

73 .  The  f  errocyanid  ion  gives  a  precipitate  of  reddish-brown 
cupric  ferrocyanid,  Cu2Fe(CN)e,  insoluble  in  dilute  acids. 

74.  Metallic   iron   or  zinc   precipitates   reddish   metallic 
copper  which  collects  on  the  metal. 

75.  Cupric  chlorid  gives  a  light  blue  color  to  the  Bunsen 
flame;  most  other  copper  compounds  give  an  emerald-green. 

Bismuth,  Bi 

Bismuth  will  dissolve  in  nitric  acid,  forming  the  nitrate, 
and  is  not  easily  attacked  by  other  acids.  It  is  a  weak  base 
and  its  salts  readily  undergo  hydrolysis,  the  bismuth  being 
precipitated  as  a  basic  salt,  and  most  of  the  acid  being  set  free. 
Hence  its  solutions  have  an  acid  reaction,  and  free  acid  must 
be  present  to  keep  it  dissolved. 

Use  for  the  reactions  a  2-per-cent.  solution  of  Bi(N03)3. 


60  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

76.  The  sulphid  ion  precipitates  black  bismuth  sulphid, 
Bi2S3,  soluble  in  nitric  acid. 

77.  The  hydroxid  ion  precipitates  Bi(OH)3,  a  white  bulky 
solid,  insoluble  in  excess. 

78.  A  bismuth  solution,  free  from  a  large  amount  of  nitric 
acid,  after  dilution  with  ten  to  twenty  times  its  volume  of 
water  undergoes  hydrolysis  and  a  white  precipitate  is  formed. 
Instead  of  the  hydroxid  which  would  be  expected,  the  reaction 
is  reversible  and  a  basic  salt  is  produced.     Boiling  or  the  addi- 
tion of  a  few  drops  of  hydrochloric  acid  favors  the  precipitation. 

Bi(N03)  3+H20  =  BiON03+2HN03. 

79.  The  stannous  ion  added  to  a  solution  of  bismuth  pre- 
viously made  alkaline  by  sodium  hydroxid  reduces  the  bis- 
muth compound  to  the  black  oxid,  Bi2O2.     This  test  can  also 
be  applied  to  many  of  the  solid  compounds. 

2Bi(NO3)3+SnCl2+ioNaOH=Bi202+6NaN03+ 
2NaCH-Na2SnO3+5H20. 

80.  The  dichromate  ion  produces  a  yellow  precipitate  of 
bismuth  oxychromate,  (BiO)2Cr04,  which  is  soluble  in  nitric 
acid  and  insoluble  in  sodium  hydroxid. 

8 1 .  The  carbonate  ion  precipitates  whi  te  basic  bismuth  car- 
bonate, or  subcarbonate. 

2Bi(N03) +3Na2CO3+H20  =  (BiO)2CO3.H2O 

+6NaN03+2C02,  or 
2Bi+3C02//+H20=(BiO)2C03.H20+2C02. 

Cadmium,  Cd 

Cadmium  dissolves  in  nitric,  hydrochloric  or  dilute  sul- 
phuric acid  with  the  formation  of  the  corresponding  salts. 

A  2-per-cent.  solution  of  Cd(NO3)2  may  be  used  for  the 
reactions. 

82.  The  sulphid  ion  precipitates  from  cadmium  solutions 


METALS  (CATIONS)  61 

bright-yellow  cadmium  sulphid,  CdS,  which  dissolves  in  nitric 
acid  but  not  in  solutions  of  potassium  cyanid  or  alkaline 
sulphids. 

83.  The  hydroxid  ion  gives  a  precipitate  of  white  cadmium 
hydroxid,  Cd(OH)2,  which  is  soluble  in  an  excess  of  am- 
monium hydroxid. 

84.  The  carbonate  ion  precipitates  white  cadmium  car- 
bonate, insoluble  in  excess. 

Mercury 

Mercury  cannot  be  dissolved  in  hydrochloric  acid.  In 
cold,  dilute  nitric  acid  it  dissolves  to  mercurous  nitrate, 
HgNOs;  in  hot,  concentrated  nitric  acid  it  also  dissolves, 
forming  then  mercuric  nitrate,  Hg(NO3)2.  Salts  of  mercury 
volatilize  upon  heating  to  a  high  temperature  in  the  absence 
of  water  either  without  or  with  decomposition.  All  the 
soluble  compounds  are  poisonous. 

The  Mercuric  Ion,  Hg 

For  the  reactions  a  2-per-cent.  solution  of  Hg(NO3)2  or 
HgCl2  may  be  employed. 

85.  The  sulphid  ion  with  mercuric  solutions  forms  a  white 
precipitate,  if  a  very  small  amount  of  the  reagent  is  used.     As 
the  reagent  is  increased  the  color  changes  to  yellow;  then 
brown  and,  finally,  black,  insoluble  mercuric  sulphid,  HgS,  is 
formed.     This  is  insoluble  in  yellow  ammonium  sulphid,  also 
in'  nitric  or  hydrochloric  acids  when  used  singly.     It  does, 
however,  dissolve  easily  in  aqua  regia. 

86.  The  hydroxid  ion,  if  obtained  from  sodium  hydroxid  or 
potassium  hydroxid,  produces  a  yellow  precipitate  of  mercuric 
oxid,  HgO. 

If  ammonium  hydroxid  is  used  as  the  reagent  the  NH3 
present  unites  with  the  mercuric  ion  to  form  a  complex  ion, 


62  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Hg+NH3  =  H'  -f  HgNHV  .     This  with  Cl'  precipitates  white 
mercuric  aminochlorid,  NH2HgCl. 

HgCl2+2NH40H-NH2HgCl+NH4CH-2H20. 

87.  The  addition  of  a  slight  amount  of  the  stannous  ion  in 
the  presence  of  HC1  forms  with  the  mercuric  ion  a  white 
precipitate  of  mercurous  chlorid,  HgCl. 

',  or 


If  more  of  the  stannous  ion  is  added  the  white  precipitate  is 
changed  to  gray  metallic  mercury  which  only  very  slowly 
collects  into  large  globules. 


Other  reducing  agents  produce  HgCl,  or,  if  strong,  gray  mer- 
cury in  a  similar  manner. 

88.  A  strip  of  sheet  copper  in  a  solution  of  mercury 
acidified  with  hydrochloric  acid  (but  free  from  nitric  acid) 
precipitates  the  mercury  as  a  gray  film  of  the  metal  on  the 
surface  of  the  copper. 


By  rubbing  with  a  cloth  or  the  finger  this  becomes  a  bright, 
white  copper  amalgam.  If  the  strip  is  inserted  in  a  narrow 
test-tube  and  this  is  heated  the  mercury  is  volatilized  and 
collects  as  a  gray  ring  in  the  cool  part  of  the  tube.  When 
examined  with  the  microscope  it  is  seen  to  be  composed  of 
globules  of  the  metal.  When  these  are  rubbed  with  a  glass 
rod  they  collect  into  larger  globules.  If  a  single  crystal  of 
iodin  is  dropped  into  the  tube  and  this  is  then  gently  warmed 
the  metal  is  converted  into  the  bright  scarlet  iodid,  HgI2. 
Zinc  and  iron,  like  copper,  convert  the  mercury  from  the  ionic 
to  the  elementary  form. 

89.  The  iodid  ion  with  mercuric  solutions  forms  a  red  pre- 


METALS  (CATIONS)  63 

cipitate  of  mercuric  iodid,  yellow  at  first  but  soon  becoming 
scarlet.  It  is  soluble  in  excess  of  either  potassium  iodid  or 
mercuric  chlorid. 

90.  Dry  compounds  of  mercury  if  mixed  with  dry  sodium 
carbonate  and  heated  in  a  tube  are  decomposed,  the  mercury 
appearing  in  globules  on  the  upper  part  of  the  glass. 

91.  Besides  the  above  reagents  many  animal  and  vegetable 
substances  form  insoluble  compounds  with  mercuric  solutions, 
e.g.,    albumin,    alkaloids,    extractive   matter,    and   a1  great 
number  of  others. 

Division  B 

The  sulphids  of  this  division  unite  with  sulphids  of  the 
alkalies  to  form  soluble  compounds  which  are  called  sulpho- 
or  thio-salts.  These  can  be  decomposed  by  acids,  hydrogen 
sulphid  being  set  free  and  the  metals  reprecipitated  as  sulphids. 

Arsenic,  As 

Metallic  arsenic  is  oxidized  by  dilute  nitric  acid  to  arsenous 
acid,  and  by  concentrated  nitric  acid  to  arsenic  acid,  both  of 
which  are  soluble.  It  is  insoluble  in  hydrochloric  and  dilute 
sulphuric  acids.  Almost  all  its  compounds  are  poisonous. 

Dry  Reactions  of  Arsenic  and  its  Compounds 

92.  Metallic  arsenic  volatilizes  when  heated,  giving  a  garlic 
odor.     If  the  heating  is  done  in  a  glass  tube  the  metal  is  depos- 
ited upon  the  upper,  cool  part  of  the  tube  as  a  coating  with  a 
metallic  luster.     When  in  a  rather  thick  layer  it  appears  black 
by  transmitted  light,  but  it  is  of  a  brown  shade  when  only  a 
minute  amount  is  present. 

93.  If  a  small  fragment  of  the  metal  is  placed  in  the  middle 
of  a  narrow  glass  tube  a  few  inches  in  length  and  open  at  both 
ends  and  is  then  heated  over  the  Bunsen  flame,  holding  the 


64  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

tube  obliquely,  the  arsenic  is  oxidized  and  the  white  crystal- 
line As20s  is  deposited  in  the  upper  part. 

94.  Arsenous  oxid  As2O3,  if  heated  in  a  glass  tube,  sub- 
limes without  an  odor,  and  is  deposited  upon  the  cooler  part 
of  the  tube  in  octahedral  or  tetrahedral  (eight-sided  or  four- 
sided)  crystals.     Their  shape  can  be  plainly  seen  with  a 
microscope.     The  best  crystals  are  obtained  by  having  the 
upper  part  of  the  tube  slightly  warmed  so  as  to  prevent  too 
rapid  a  deposition  of  the  oxid. 

95.  Arsenous  oxid  when  volatilized  over  glowing  charcoal 
is  reduced  to  the  metal.     This  can  be  demonstrated  by  draw- 
ing out  the  closed  end  of  a  small  tube  to  a  rather  narrow  point 
and,  after  it  has  cooled,  placing  a  fragment  of  the  oxid  in  the 
end,  with  a  splinter  of  charcoal  above  it.     Heat  the  charcoal 
to  redness  first,  then  gradually  heat  the  arsenic  compound. 
A  dark  mirror  of  arsenic  appears  in  the  upper  part  of  the  tube. 
Some  other  compounds  of  arsenic  will  give  the  same  result, 
but  not  all  of  them. 

96.  Arsenic  sulphid  and  some  of  the  arsenic  compounds 
which  are  not  reduced  to  the  metallic  form  by  charcoal  alone, 
will  be  thus  changed  by  mixing  with  sodium  carbonate  and 
potassium  cyanid,  both  previously  thoroughly  dried,  and 
heating  in  the  tube.     The  mirror  appears  as  before. 

Reactions  of  Arsenic  Ions 

In  making  the  reactions  of  the  ions  of  arsenic  we  may  con- 
sider them  as  belonging  to  two  classes,  first  those  in  which  the 
arsenic  acts  as  the  cation,  as  it  does  in  the  halogen  compounds, 
and,  secondly,  the  arsenous  and  arsenic  acids  and  their  salts 
(arsenites  and  arsenates)  where  the  arsenic  is  found  in  the 
negative  part  or  anion.  Some  reactions  will  be  common 
to  all  compounds;  others  will  give  different  results  with  the 
different  classes. 


METALS  (CATIONS)  65 

97.  Arsenous'  oxid  is  soluble  in  water,  though  with  diffi- 
culty.    When  so  dissolved  it  forms  the  weak  acid,  H3AsOs, 
arsenous  acid.     It  dissolves  in  hydrochloric  acid  to  arsenous 
chlorid,    AsCl3.     Sodium  hydroxid  or  potassium  hydroxid 
dissolves  arsenous  oxid  easily  with  the  formation  of  the  arsen- 
ites  of  these  metals.     Nitric  acid  and  also  aqua  regia  convert 
it  into  soluble  arsenic  acid,  H3AsO4. 

General  Reactions  of  the  Arsenous  Ion 

Use   a  i-per-cent.  solution  of  AsCl3,  As2O3,  or  K3As03* 

98.  The  sulphid  ion  with  neutral  or  alkaline  solutions 
gives  no  precipitate,  but  upon  the  addition  of  hydrochloric 
acid  it  precipitates  the  arsenic  as  bright  yellow  arsenous 
sulphid,  As2S3.     This  is  soluble  in  the  alkaline  hydroxids, 
carbonates,  or  sulphids,  but  not  in  hydrochloric  acid.     When 
it  is  dissolved  in  ammonium  sulphid  ammonium  thioarsenite, 
or  sulpharseriite  is  formed,  NH4AsS2.     This  is  decomposed  by 
acids,  As2S3  being  reprecipitated.     Boiling  nitric  acid  dissolves 
it  after  converting  it  into  arsenic  acid. 

Reactions  of  the  Arsenite  Ion  (Arsenous  Acid  and  Arsenites) 

Use  a  i-per-cent.  solution  of  As2O3  or  K3AsOs. 

99.  The  silver  ion  in  neutral  solutions  of  an  arsenite  iorms 
yellow  silver  arsenite,  Ag3AsO3,  soluble  both  in  excess  of  am- 
monia and  of  nitric  acid.     This  reaction  :'s  concealed  by  the 
presence  of  any  acid  or  salt  which  precipitates  silver  nitrate, 
such  as  hydrochloric  acid  or  the  chlorids.     A  convenient  way 
to  make  the  reaction  is  to  first  acidify  the  arsenous  solution 
with  nitric  acid  in  a  narrow  test-tube,  then,  after  adding  a  few 
drops  of  silver  nitrate,  to  slant  the  tube  and  pour  ammonium 
hydroxid  down  the  side  carefully  so  as  to  prevent  the  liquids 
from  mixing.     The  precipitate  appears  in  the  neutral  zone 
between  them. 


66  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

TOO.  The  copper  ion  in  neutral  solutions  in  the  same  way 
produces  yellowish-green  cupric  arsenite,  CuHAs03.  It  dis- 
solves in  either  ammonium  hydroxid  or  nitric  acid. 

101.  Cations  of  most  metals  except  those  of  the  alkalies 
form  insoluble  arsenites  in  neutral  solutions. 

General  Reactions  of  the  Arsenic  Ion 


Use  a  i-per-cent.  solution  of  N 

1  02.  The  sulphid  ion  gives  no  precipitate  in  neutral  or  alka- 
line solutions  of  arsenates.  If  they  are  acidified  it  slowly  re- 
duces them  to  arsenous  compounds,  then  precipitates  them  as 
arsenous  sulphid,  As2Sa.  Warming  hastens  the  reduction 
and  precipitation,  but  several  hours  are  required  for  its  com- 
pletion —  sometimes  twenty-four.  Sulphur  is  precipitated  at 
the  same  time. 

103.  With  the  sulphid  ion  in  an  alkaline  solution,  or  if 
ammonium   sulphid   is   used,   no   precipitate   is   produced, 
although  the  soluble  ammonium  thioarsenate,  (NH^sAsS^  is 
formed.     When  this  is  acidified  with  hydrochloric  acid  yellow 
arsenic  sulphid,  As2S5,  is  precipitated. 

Reactions  of  the  Arsenate  Ion  (Arsenic  Acid  and  Arsenates) 

Use  a  i-per-cent.  solution  of  Na2HAsC>4. 

104.  The  silver  ion  produces  in  neutral  solutions  of  arsen- 
ates a  reddish-brown  precipitate  of  silver  arsenate,  Ag3AsO4, 
soluble  in  dilute  nitric  acid  and  ammonium  hydroxid.     The 
reaction  may  be  made  in  the  same  manner  as  in  99. 

105.  The  copper  ion  precipitates  from  neutral  solutions 
greenish-blue   cupric   arsenate,    CuHAsO4,   soluble   in   am- 
monium hydroxid  or  nitric  acid. 

1  06.  Magnesia  mixture  (magnesium  sulphate  with  am- 
monium chlorid  and  ammonium  hydroxid;  it  must  be  a  clear 
solution  and  not  cloudy)  forms  white  crystalline  ammonium 


METALS  (CATIONS)  <  67 

magnesium  arsenate  with  the  arsenate  ion.     It  dissolves 
easily  in  acids. 

107.  Cations  of  most  metals  except  the  alkali  metals  form, 
in  neutral  liquids,  insoluble  arsenates. 

General  Reactions  of  Arsenous  and  Arsenic  Ions 

Use  a  i-per-cent.  solution  of  any  compound  of  arsenic. 

1 08.  From   solutions  of  arsenic,   strongly  acidified  with 
hydrochloric  acid  and  warmed  with  strips  of  sheet  copper, 
copper  arsenid  is  precipitated  which  forms  a  dark-gray  film 
on  the  surface  of  the  copper.     With  very  dilute  solutions  this 
can  be  accomplished  by  placing  the  solution  with  the  copper 
in  a  test-tube  and  letting  it  stand  half  an  hour  in  a  beaker  of 
boiling  water.     If  then  the  copper  is  removed,  washed  with 
water,  dried  and  heated  in  a  small  open  tube  (93)  the  arsenic 
is   oxidized   to   As2O3,   which  is   deposited  in  microscopic 
crystals.     The  copper  remains  black  from  the  formation  of  its 
oxid.     This  test  (Reinschs'  test)  is  very  sensitive  and  is  not 
prevented  by  the  presence  of  organic  matter. 

109.  Acidify  3-4  c.c.  of  arsenic  solution  with  pure  hydro- 
chloric or  sulphuric  acid,  add  a  small  piece  of  zinc  (about  one 
gram)  that  is  known  to  be  free  from  arsenic,  cover  the  mouth 
of  the  tube  with  several  thicknesses  of  filter-paper  upon  the 
upper  one  of  which  has  been  put  a  drop  of  saturated  silver 
solution.    Let  the  tube  stand  where  it  is  protected  from  a 
bright   light   and   from   hydrogen   sulphid.     After   a   time, 
varying  with  the  amount  of  arsenic  present,  the  spot  on  the 
paper  becomes  yellow  with  a  brown  or  black  margin.     The 
addition  of  water  changes  it  to  black.     The  yellow  com- 
pound is  Ag3As,  3AgN03.     This  is  produced  by  the  action  of 
the  AsH3 — formed  by  the  nascent  hydrogen  with  the  soluble 
arsenic  compound — and  the  silver  nitrate.     Water  decom- 
poses it,  setting  free  metallic  silver. 


68 


INTRODUCTION  TO   CHEMICAL  ANALYSIS 


Gutzeit's  test. 

Ag3AS,3AgN03+3H20  =  6Ag+H3AsO3-f  3HNO3. 
A  mercuric  chlorid  solution  can  be  substituted  for  the  sat- 
urated silver  nitrate  and  gives  a  yellow  color  also. 
This  test  is  a  very  sensitive  one. 

no.  If  sodium  hydroxid  or  potassium  hydroxid  is  substi- 
tuted for  the  acid,  the  reaction  being  otherwise  performed  in 
the  same  way,  except  that  the  solution  may  be  slightly 
warmed  the  same  result  appears  (Fleit- 
mann's  test).  This  reaction  is  of  value 
in  distinguishing  between  arsenic  and 
antimony. 

in.  To  3-4  c.c.  of  an  arsenic  solution 
add  an  equal  volume  of  concentrated 
hydrochloric  acid  and  half  a  gram  of 
metallic  tin,  granulated  or  in  the  form 
of  foil.  Upon  warming  for  some  time 
a  brown  color  or  precipitate  of  arsenic 
is  produced  (Bettendorf's  test). 

112.  Nascent  hydrogen  converts  ar- 
senic compounds  into  gaseous  hydrogen 
arsenic,  or  arsine,  AsH3.  This  can  be 
decomposed  either  by  heating  or  by  cool- 
ing the  flame  of  the  burning  gas,  the 
arsenic  being  deposited  as  a  dark  mirror- 
like  coating  on  cold  objects.  This  is  the 
basis  of  Marsh's  test.  For  making  this 
many  forms  of  apparatus  have  been 
proposed.  The  simplest  of  these  consist  of  a  small  flask 
(200-300  c.c.)  fitted  with  a  doubly  perforated  stopper,  one  of 
rubber  being  preferable  to  cork.  Through  one  hole  a  funnel- 
tube  (a  funnel  with  a  piece  of  rubber  tubing  slipped  over  the 
stem  answers  for  this)  passes  nearly  to  the  bottom  of  the  flask. 
The  other  is  provided  with  an  exit  tube  of  difficultly  fusible 


FIG.  ii. — A  simple 
form  of  generator  for 
Marsh's  test. 


METALS  (CATIONS)  69 

glass  which  does  not  project  into  the  flask  and  of  which  the 
outer  end  is  drawn  out  to  a  small  point.  Into  this  apparatus 
is  put  5-10  grams  of  pure  granulated  zinc,  the  stopper  is 
inserted  and  dilute  sulphuric  (one  to  four),  is  poured  in 
through  the  funnel  tube.  The  hydrogen  is  allowed  to  escape 
until  it  is  judged  that  all  the  air  has  been  expelled  before  the 
gas  is  ignited.  If  there  is  any  doubt  of  this  a  towel  should 
first  be  loosely  wrapped  about  the  flask  to  avoid  any  flying 
fragments  of  glass  if  there  is  an  explosion.  Notice  the  color 
of  the  hydrogen  flame,  then  hold  in  it  a  cold  porcelain  dish, 
moving  it  about  so  as  to  prevent  its  becoming  excessively 
heated.  If  the  zinc  and  acid  are  both  free  from  arsenic  there 
will  be  no  discoloration  of  the  porcelain.  Now  pour  into  the 
hydrogen  flask  a  few  drops  of  the  solution  to  be  tested  for 
arsenic,  avoiding  as  far  as  possible  the  introduction  of  air  with 
it.  The  color  of  the  flame  soon  changes  to  a  bluish- white  and 
upon  the  dish  there  will  be  deposited  a  brownish  or  black 
mirror  of  arsenic.  The  gas  should  not  be  allowed  to  escape 
into  the  room  unburned  after  the  arsenic  compound  is 
present  since  the  hydrogen  arsenid  is  extremely  poisonous. 
Apply  the  following  tests  to  the  arsenic  mirror. 

(a)  Drop  upon  it  a  little  of  a  solution  of  chlorinated  lime 
(calcium  hypochlorite) .  The  mirror  dissolves. 

(6)  Add  a  drop  of  concentrated  nitric  acid.  This  also  dis- 
solves the  arsenic.  Carefully  evaporate  this  solution  to  dry- 
ness,  without  heating  to  a  high  temperature.  Moisten  it 
with  a  drop  of  silver  nitrate.  A  brick-red  color  will  be  pro- 
duced if  all  acid  has  been  removed. 

(c)  Add  to  the  mirror  a  few  drops  of  ammonium  sulphid. 
The  arsenic  is  dissolved  and  the  solution,  if  evaporated  to 
dryness,  leaves  a  yellow  residue  of  arsenic  sulphid. 

Where  the  test  is  to  be  made  more  carefully  for  minute 
amounts  of  arsenic,  a  somewhat  larger  generating  flask  is 
used  (perhaps  one  liter)  with  more  of  the  zinc.  The  funnel 


70  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

tube  should  be  provided  with  a  stop-cock  to  prevent  the  ad- 
mission of  air.  The  gas  will  contain  moisture  from  the  di- 
lute acid,  and  to  remove  this  should  be  passed  first  through  an 
empty  U-tube  cooled  by  suspension  in  water,  then  through  a 
tube  which  is  filled  with  lumps  of  dried  calcium  chlorid.  The 
latter  absorbs  the  moisture  which  is  not  condensed  in  the  first 
tube.  The  dry  gas  now  passes  through  a  horizontal  tube  of 
glass  which  does  not  soften  except  at  a  very  high  temperature 
("hard"  or  "combustion"  glass).  This  tube  is  to  be  heated 


FIG.  12. — Apparatus  for  the  detection  of  a  minute  amount  of  arsenic  by 

Marsh's  test. 


to  redness  with  a  Bunsen  burner  after  all  air  has  been  driven 
out  of  the  apparatus.  Before  introducing  the  arsenic  solution, 
while  the  hydrogen  is  passing  through  it,  the  tube  should  be 
heated  for  half  an  hour  to  prove  the  purity  of  the  reagents 
used.  The  presence  of  arsenic  is  denoted  by  a  brown  or 
black  mirror  in  the  part  of  the  tube  immediately  beyond  the 
flame.  If  the  glass  is  at  this  place  contracted  to  a  small  diam- 
eter the  mirror  is  more  evident.  If  any  of  the  arsenic  com- 
pound is  not  decomposed  here,  which  is  usually  the  case,  it  can 
be  tested  for  in  the  flame  at  the  pointed  end  of  the  tube  by 
means  of  the  cold  porcelain,  -as  described  above.  During  the 
test  the  escaping  gas  should  be  burned  at  the  end  of  the  tube 


METALS  (CATIONS)  71 

passed  into  a  solution  of  silver  nitrate  to  prevent  the  hydro- 
gen arsenid  from  escaping  into  the  air.  With  silver  nitrate 
it  forms  a  precipitate  of  metallic  silver  (black)  and  the  arsenic 
dissolves  in  the  liquid.  In  long  operations  it  is  advisable  to 
provide  a  means  of  removing  the  exhausted  acid  without  the 
admission  of  air.  This  may  be  affected  by  inserting  through 
a  third  hole  in  the  stopper  a  bent  glass  tube  closed  at  the  outer 
end  by  a  clamped  rubber  tube.  By  opening  the  clamp  and 
closing  the  exit  tube  for  a  second  the  liquid  can  be  siphoned 
from  the  flask.  Where  very  minute  amounts  of  arsenic  are 
sought  the  operation  may  last  for  several  hours. 

The  arsenic  mirror  in  the  tube,  when  gently  warmed  in  a 
slow  stream  of  dry  hydrogen  sulphid,  is  converted  into  yellow 
arsenous  sulphid.  The  reactions  made  with  the  mirror  on 
porcelain  can  be  applied  here  also. 

By  Marsh's  test  less  than  one  one-hundredth  of  a  milli- 
gram of  arsenic  in  a  cubic  centimeter  has  been  detected.  It 
is  interfered  with  or  prevented  by  the  presence  of  oxidizing 
agents,  organic  matter,  and  the  salts  of  some  heavy  metals 
like  mercury. 

113.  Freshly  precipitated  ferric  hydroxid   (54)    removes 
arsenous  acid  or  arsenites  from  their  solutions,  as  can  be  seen 
by  shaking  an  excess  with  the  arsenical  liquid,  filtering  and 
testing  the  filtrate  for  the  arsenic.     Magnesium  hydroxid  or 
oxid  acts  in  a  similar  manner. 

114.  The  hydroxid  ion  and  the  carbonate  ion  precipitate 
neither  arsenous  nor  arsenic  ions. 


Antimony,  Sb 

The  metal  is  of  a  white  color  with  a  bright  luster  and  is  very 
brittle.  It  does  not  dissolve  in  hydrochloric  acid,  but  is  con- 
verted to  a  chlorid  by  aqua  regia.  Concentrated  nitric  acid 
changes  it  into  a  white  metantimonic  acid  that  is  nearly  insol- 


72  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

uble  in  the  nitric  acid.     Its  soluble  compounds,  in  large  doses, 
are  poisonous. 

For  the  reactions  a  i-per-cent.  solution  of  KSbC^H^Oe 
may  be  used. 

115.  The  sulphid  ion  precipitates  from  acid  solutions  orange- 
red  antimonous  sulphid,  Sb2Sa.     It  is  soluble  in  ammonium 
sulphid,  as  well  as  in  the  other  alkaline  sulphids,  to  the  sul- 
phantimonate  which,  like  the  corresponding  arsenic  compound, 
is  decomposed  by  acids  with  the  precipitation  of  Sb2S5.     It  is 
insoluble  in  ammonium  carbonate.     Boiling  concentrated  hy- 
drochloric acid  dissolves  it,  but  it  is  insoluble  in  dilute  acids. 

116.  Metallic  zinc  precipitates  from  solutions  which  con- 
tain no  nitric  acid  fine,  black,  metallic  antimony.     If,  in  an 
antimony  solution  containing  an  excess  of  hydrochloric  acid, 
a  piece  of  platinum  foil  is  brought  in  contact  with  the  zinc,  the 
antimony  is  deposited  as  a  dark-brown  or  black  stain  on  the 
platinum.     This  can  be  removed  by  hot  nitric  acid,  but  does 
not  dissolve  in  hydrochloric  acid. 

117.  With   nascent   hydrogen,   in   a   Marsh's   apparatus 
(112),  the  antimony  is  converted  into  hydrogen  antimonid, 
SbH3,  a  gas  which  is  decomposed  like  the  AsH3  with  the 
formation  of  an  antimony  mirror.     It  has  these  points  of 
difference : 

(a)  It  is  usually  of  a  more  sooty  black,  although  this  may 
not  be  evident  where  only  a  very  small  amount  is  present. 

(b)  In  the  tube  of  the  Marsh's  apparatus  it  is  formed  on 
both  sides  of  the  spot  where  the  glass  is  heated,  and  is  not  so 
easily  volatile  as  the  arsenic. 

(c)  It  does  not  dissolve  in  a  solution  of  calcium  hypochlorite. 

(d)  It  dissolves  in  a  drop  of  concentrated  nitric  acid,  but 
upon  subsequent  evaporation  to  dryness  and  moistening  with 
silver  nitrate  solution  it  remains  colorless. 

(e)  When  warmed  in  a  current  of  dry  hydrogen  sulphid 
gas  it  is  changed  into  red  or  black  antimonous  sulphid. 


METALS  (CATIONS)  73 

(/)  In  ammonium  sulphid  it  dissolves  and  upon  evapora- 
tion to  dryness  the  orange  red  sulphid,  Sl^Sa,  remains. 

1 18.  Many  of  the  soluble  antimony  salts  are  hydrolyzed  by 
large  quantities  of  water  with  the  precipitation  of  basic  salts. 
This  is  true  of  the  chlorid  which  is  changed  into  the  basic 
chlorid.     It  sometimes  occurs  also  when  antimony  solutions 
are  slightly  acidified  by  hydrochloric  acid.     The  precipitate 
dissolves  in  a  larger  amount  of  the  acid  or  in  tartaric  acid  and 
is  prevented  from  appearing  if  the  tartaric  acid  is  added 
previously. 

Sb-  +  3C1'  +  H20  =  SbOCl  +  2H'  +  2C1'. 

119.  With  Remsch's  test  (108)  antimony  compounds  give 
a  dark  coating  on  the  copper,  but  by  heating  in  the  open  tube 
this  forms  amorphous  Sb2Oa  which  can  be  distinguished  from 
the  arsenic  compound  by  the  aid  of  the  microscope  (94). 

120.  Gutzeit's  test  (109)  gives  immediately  a  black  or 
brown  stain  of  silver  antimonid,  Ag3Sb,  on  the  paper,  not  a 
yellow  one. 

121.  Fleitmann's  test  (no)  gives  no  result  with  antimony 
compounds. 

122.  Tannic  acid  with  the  antimony  ion  forms  a  white  pre- 
cipitate of  antimony  tannate,  soluble  in  tartaric  acid. 

123.  WithBettendorf's  test  (in)  the  antimony  compounds 
are  decomposed,  antimony  being  deposited  as  a  black  coating 
on  the  tin. 

124.  The  hydroxid  ion  or  the  carbonate  ion  precipitates 
white  antimonous  hydroxid  Sb(OH)3. 

Tin,  Sn 

The  metal  dissolves  in  concentrated  hydrochloric  acid, 
forming  stannous  chlorid.  Concentrated  nitric  acid  changes 
it  into  a  white  metastannic  acid  insoluble  in  excess.  In  con- 
centrated aqua  regia  tin  dissolves  to  stannic  chlorid. 

There  are  two  series  of  tin  compounds,  the  stannous  and 


74  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

the  stannic.     The  former  are  readily  converted  into  the  latter 
by  oxidizing  agents  or  by  standing  exposed  to  the  air. 

The  Stannous  Ion  (Sn) 

Use  a  2-per-cent.  solution  of  SnCl2  for  the  reactions. 

125.  The  sulphid  ion  in  neutral  or  acid  solutions  precipi- 
tates stannous  sulphid,  SnS,  a  dark  brown  compound.     It 
dissolves  in  boiling  hydrochloric  acid.     It  is  insoluble  in 
colorless  ammonium  sulphid,  but  soluble  in  the  yellow  ammo- 
nium sulphid  to  a  sulpho-stannate,  (NH4)2SnS3.     From  this 
solution  acids  precipitate  yellow  stannic  sulphid,  SnS2,  mixed 
with  sulphur. 

126.  The  hydroxid  ion  precipitates  white  stannous  hy- 
droxid,  Sn(OH)2.     It  is  insoluble  in  excess  of  ammonium 
hydroxid,  but  is  soluble  in  excess  of  sodium  hydroxid  or 
potassium  hydroxid,   forming  Na2SnO2,  from  which  com- 
pound boiling  precipitates  stannous  oxid,  SnO. 

127.  Metallic  zinc  precipitates  dark  gray  or  black,  spongy, 
metallic  tin  from  solutions  containing  hydrochloric  acid.     If 
platinum  foil  is  present  it  is  not  discolored. 

128.  The  mercuric  ion  in  hydrochloric  acid  solution  is  re- 
duced to  the  mercurous  ion  and  (compare  87)  gives  a  white 
precipitate  of  mercurous  chlorid,  HgCl,  which  turns  black 
upon  being  made  alkaline  with  ammonium  hydroxid. 

The  Stannic  Ion  (Snj 

Use  for  reactions  a  2-per-cent.  solution  of  SnCU. 

129.  The  sulphid  ion  in  excess  precipitates  from  acid  or 
neutral  solutions  yellow  stannic  sulphid,  SnS2.     It  dissolves 
in  boiling  hydrochloric  acid  and  in  the  alkaline  sulphids  to  a 
sulpho-salt  (125)  and  is  reprecipitated  unchanged  by  acids. 

130.  Zinc  gives  the  same  result  as  with  the  stannous  ion. 


METALS  (CATIONS)  75 

Gold,  Au 

Gold  is  insoluble  in  nitric,  hydrochloric,  or  sulphuric  acids. 
It  dissolves  in  aqua  regia  or  in  other  liquids  where  chlorin  is 
set  free,  forming  the  chlorid  AuCl3,  and  also  is  soluble  in 
potassium  cyanid. 

A  very  dilute  solution  of  the  chlorid  may  be  used  for  the 
reactions. 

131.  The  sulphid  ion  precipitates  solutions  of  gold  <s  a 
dark  brown  auric  sulphid,  Au2Ss,  that  is  soluble  in  alkaline 
polysulphids  and  in  aqua  regia,  but  not  in  single  acids. 

132.  The  ferrous  ion  precipitates  from  acid  solutions  metal- 
lic gold  as  a  yellowish-brown  powder. 

2Au  +  6Fe  =  2Au  +  6Fe,  or 

2AuCl3  +  6FeS04  +  3H2SO4  =  Au2  +  3Fe2(SO4)3  +6HC1. 

133.  Oxalic  acid,  H2C2O4,  also  precipitates  the  metallic 
gold  from  its  solutions. 

2Au"+  3C2(V'  =  2Au  +  6CO2,  or 
2AuCl3  +  3H2C204  =  Au2  +  6HC1  +  6CO2. 

The  last  two  reagents  precipitate  the  gold  because  of  their 
power  of  reduction.  Other  reducing  agents  act  in  a  similar 
manner. 

134.  Stannous  chlorid  containing  a  little  stannic  chlorid, 
SnCl4,  formed  by  partly  oxidizing  the  stannous  solution, 
forms  with  gold  solutions  a  purple  compound  of  varying 
composition,  called  the  Purple  of  Cassius.     It  remains  in 
suspension  giving  a  purple  color  to  the  .liquid. 

135.  Metallic  zinc,  iron,  copper  or  mercury,  as  well  as  some 
other  metals,  will  precipitate  the  gold  in  the  dark,  finely 
divided  metallic  form. 

136.  Ammonium  hydroxid,  in  concentrated  gold  solutions, 
forms  a  precipitate  of  reddish-brown  fulminating  gold  which, 
when  dry,  explodes  violently  by  friction. 


7  6  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

137.  Compounds  of  gold  are  decomposed,  the  metal  being 
precipitated,  by  organic  matter,  e.g.,  by  boiling  with  sugar. 

138.  Very  minute  quantities  of  gold  can  be  detected  by  ex- 
tracting the  metal  from  the  finely  powdered  material  by 
means  of  tincture  of  iodin.     It  shoud  be  left  in  contact  for 
some  time  to  insure  an  effective  solution.     If,  then,  a  piece  of 
absorbent  white  paper  is  dipped  in  the  solution  and  afterward 
dried  and  burned  the  gold  will  color  the  ash  purple.     The  test 
is  more  delicate  if  the  paper  is  dipped  and  dried  several  times 
before  ignition. 

Platinum,  Pt 

The  metal  is  not  soluble  in  single  acids,  but  .it  dissolves  in 
aqua  regia,  although  not  so  readily  as  does  gold.  This  solu- 
tion contains  hydrochlorplatinic  acid,  H2PtCl6. 

A  i-per-cent.  solution  of  H2PtCle  may  be  used  for  the 
reactions. 

139.  The  sulphid  ion  in  acid  solutions  precipitates  slowly 
(best  with  the  aid  of  heat)  platinic  sulphid,  PtS2.     This  is  a 
nearly  black  compound,  insoluble  in  single  acids,  but  soluble 
in  alkaline  polysulphids.     From  this,  acids  reprecipitate  the 
sulphid.     It  is  also  soluble  in  aqua  regia. 

140.  The  ammonium  or  potassium  ion  precipitates  from 
hydrochloric  acid  solutions  a  double  salt — ammonium  chlor- 
platinate,  (NH^PtCle,  or  potassium  chlorplatinate,  K2PtCl6. 
These  are  yellow  crystalline  salts,  somewhat  soluble  in  water 
and  rendered  less  so  when  alcohol  is  present. 

141.  Metallic  zinc  or  iron  precipitates  the  platinum  as 
spongy  black  metallic  platinum. 

142.  Oxalic  acid  does  not  precipitate  platinum  solutions 
and  the  ferrous  ion  does  so  only  after  long  boiling.     Some  or- 
ganic substance,  however,  like  sugar  or  alcohol,  will  decom- 
pose boiling  alkaline  solutions,  the  metal  being  thrown  down 
as  a  fine  black  powder — platinum  black. 


METALS  (CATIONS)  77 

Directions  for  the  Separation  of  the  Cations  of  Group  II 

If  the  solution  does  not  contain  free  acid,  make  it  acid  by 
the  addition  of  hydrochloric.  Precipitate  all  the  ions  of  the 
group  with  hydrogen  sulphid.  This  may  be  done  by  allowing 
the  gas  to  pass  through  the  solution,  or  by  the  use  of  a  solu- 
tion of  hydrogen  sulphid  in  water  until  a  precipitate  no  longer 
appears.  Warm  and  stir  the  liquid  until  the  precipitate  set- 
tles, then  when  the  metals  have  been  completely  precipitated 
filter  and  wash.  Groups  III,  IV,  and  V  pass  into  the  filtrate. 
Group  II  is  divided  into  two  divisions,  A  and  B,  according 
to  whether  it  is  insoluble  (A)  or  soluble  (B)  in  yellow  ammo- 
nium sulphid.  Scrape  the  precipitate  from  the  filter  or  rinse 
it  off  with  the  aid  of  15  c.c.  of  yellow  ammonium  sulphid  into 
a  small  porcelain  dish.1  Heat  gently,  but  not  to  boiling  and 
keep  it  at  this  temperature  for  five  minutes,  stirring  often. 
Filter  and  wash,  discarding  the  wash  water.  If  the  precipi- 
tate has  a  tendency  to  pass  through  the  filter  this  may  be  pre- 
vented by  warming  with  5  c.c.  of  ammonium  chlorid  solution. 

Division  A— Sulphids  Insoluble  in  Yellow  Ammonium 

Sulphid 

The  precipitate,  obtained  by  saturating  the  acidified  solu- 
tion with  hydrogen  sulphid,  contains  the  sulphids  of  mercury, 
lead,  bismuth,  copper,  and  cadmium. 

After  the  precipitate  has  been  well  washed  rinse  it  from  the 
paper  with  10  c.c.  of  dilute  nitric  acid  and  boil.  The  sulphids 
are  all  dissolved  with  the  exception  of  the  black  mercuric  sul- 
phid which  remains  as  a  heavy  sediment.  A  part  of  the  lead 
sulphid  may  be  oxidized  to  lead  sulphate,  an  insoluble,  white, 

1  Instead  of  at  first  treating  the  whole  precipitate  with  ammonium  sulphid, 
it  is  better  to  test  a  small  portion  thus  to  learn  if  B  is  present.  If  it  is,  treat 
the  whole  in  the  same  manner;  if  not,  proceed  directly  to  the  examination  for 
A.  If  it  dissolves  completely  A  is  absent. 


78  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

heavy  powder.  Sulphur  may  also  be  present  as  a  grayish, 
floating  powder.  Filter  out  the  precipitate,  wash  it  and, 
after  removing  from  the  paper,  dissolve  it.  Filter  off  the 
sulphur  and  test  the  nitrate  with  a  few  drops  of  stannous 
chlorid.  The  mercury  is  precipitated  as  a  white  chlorid 
which  becomes  gray  with  an  excess  of  the  tin  solution  or  after 
heating. 

The  solution  of  lead,  bismuth,  copper,  and  cadmium  should 
be  evaporated  nearly  to  dryness  under  a  hood,  then  diluted 
to  10  c.c.  with  water  and  tested  for  lead  by  adding  dilute 
sulphuric  acid  (avoiding  a  large  excess)  and  half  a  dozen  drops 
of  alcohol.  Lead  sulphate,  a  fine,  white,  heavy  powder  is 
thrown  down.  It  should  be  allowed  to  settle  before  filtering 
and,  if  necessary,  a  little  more  alcohol  added  to  insure  a  com- 
plete precipitation.  Filter  and  wash.  The  lead  may  be  con- 
firmed by  pouring  3-4  c.c.  of  sodium  hydroxid  upon  the 
washed  lead  sulphate  without  removing  from  the  filter,  then 
adding  ammonium  sulphid  to  the  filtrate  when  the  black 
lead  sulphid  should  be  formed. 

To  the  filtrate  from  the  precipitated  lead  sulphate,  which 
contains  bismuth,  copper,  and  cadmium,  add  ammonium 
hydroxid  until  it  smells  strongly  of  this  reagent.  The  bis- 
muth is  precipitated  as  a  white  flocculent  hydroxid  while  the 
copper  and  cadmium  dissolve  in  excess,  the  former  giving  a 
deep  blue  solution.  Filter,  wash,  and  confirm  the  bismuth 
by  the  addition  of  a  few  drops  of  sodium  hydroxid  and  as 
much  stannous  chlorid,  to  the  substance  on  the  paper,  when 
the  compound  turns  black. 

If  a  blue  color  indicates  copper  in  the  filtrate  from  the  bis- 
muth, acidify  one-half  and  confirm  the  copper  by  potassium 
ferrocyanid  which  produces  a  reddish-brown  precipitate. 
To  the  remainder  add  potassium  cyanid  until  the  blue  color 
disappears,  then  precipitate  the  bright  yellow  cadmium  sul- 
phid by  hydrogen  sulphid.  If  no  blue  color  is  present  the 


METALS  (CATIONS)  79 

copper  is  probably  absent,  in  which  case  the  potassium  cyanid 
may  be  omitted  and  the  whole  filtrate  from  the  bismuth  may 
be  tested  for  cadmium. 


Division  B — Sulphids  Soluble  in  Yellow  Ammonium  Sulphid 

This  contains  the  sulphids  of  arsenic,  antimony,  tin,  gold, 
and  platinum.  Acidify  the  nitrate  with  hydrochloric  acid. 
The  sulphids  of  metals  are  precipitated  in  curdy  masses.  If 
the  precipitate  is  fine  and  nearly  white  and  settles  very  slowly 
it  is  only  sulphur  and  may  be  neglected.  Otherwise  filter  and 
wash,  adding  ammonium  chlorid  as  above  if  the  filtrate  is  not 
clear.1  From  the  precipitate  which  contains  the  sulphids  of 
all  the  metals  of  Division  B,  dissolve  the  arsenic  sulphid  by 
digesting  several  minutes  with  about  10  c.c.  of  ammonium 
carbonate,  then  filter.  The  arsenic  can  be  detected  in  the  fil- 
trate by  acidifying  with  hydrochloric  acid,  then  passing  in 
hydrogen  sulphid;  yellow  arsenious  sulphid  is  precipitated. 

The  well-washed  precipitate  should  have  the  water  re- 
moved from  it  as  far  as  possible  by  draining  and  pressing, 
then  after  removal  from  the  paper  is  to  be  warmed  in  about  5 
c.c.  of  concentrated  hydrochloric  acid  for  several  minutes. 
This  should  be  done  under  a  hood  and  the  acid  must  not  be 
allowed  to  boil  violently.  The  sulphids  of  antimony  and  tin 
dissolve,  and  those  of  gold  and  platinum  remain,  mixed  with 
some  sulphur  and  possibly  a  part  of  the  arsenic  sulphid. 

After  filtering  pour  the  filtrate  into  a  dish,  add  a  few  frag- 
ments of  zinc  and  immediately  bring  a  piece  of  platinum  foil 
into  contact  with  this.  If  antimony  is  present  it  will  make  a 
brown  or  black  stain  on  the  foil  to  which  it  adheres  very 
firmly.  The  tin  is  precipitated  in  the  metallic  form,  but 
nearly  black  in  color  and  so  heavy  that  it  quickly  settles  to 

1  As  the  filtrate  is  not  to  be  used  it  may,  for  economy  of  time,  be  decanted 
without  filtration. 


80  INTRODUCTION   TO  CHEMICAL  ANALYSIS 

the  bottom.  Let  the  action  proceed  until  it  seems  certain 
that  all  the  antimony  and  tin  are  precipitated,  then  wash  the 
tin  by  decantation  and  dissolve  it  in  one  or  two  cubic  centi- 
meters of  warm  concentrated  hydrochloric  acid.  This  solu- 
tion when  added  to  mercuric  chlorid  precipitates  white 
mercurous  chlorid,  which  turns  black  if  the  liquid  is  made 
alkaline  with  ammonia.  If  the  stannous  chlorid  is  in  ex- 
cess it  sets  free  gray  metallic  mercury 'from  the  mercurous 
chlorid. 

The  antimony  can  be  confirmed,  if  thought  desirable,  by 
warming  the  coated  foil  in  2-3  c.c.  of  yellow  ammonium  sul- 
phid  in  a  test-tube,  then  evaporating  to  dryness.  Orange- 
red  antimonous  sulphid  remains. 

The  remaining  precipitate  from  which  the  sulphids  of 
arsenic,  antimony,  and  tin  have  been  dissolved  by  ammonium 
carbonate  and  hydrochloric  acid  will  be  dark  colored  if  it  con- 
tains the  sulphid  of  gold  or  platinum.  Otherwise  it  will  con- 
sist of  sulphur  with  possibly  some  arsenic  sulphid  that  the 
ammonium  carbonate  did  not  completely  dissolve.  If  it  is 
dark  in  color,  warm  it  with  aqua  regia  to  dissolve  the  gold  and 
platinum,  filtering  if  solution  is  not  complete.  Test  half  the 
filtrate  for  gold  with  a  mixture  of  stannous  and  stannic 
chlorids  (Purple  of  Cassius,  134).  From  the  other  half,  pre- 
cipitate the  yellow,  crystalline  ammonium  chlorplatinate 
with  NH4C1,  using  a  rather  concentrated  solution.  Con- 
firm this  by  examination  with  the  microscope  (140). 

TABLE  IV 

OUTLINE  OF  THE  SEPARATION  or  THE  CATIONS  or  GROUP  II 

From  the  acidified  solution  precipitate  with  H2S.  The  precipitate  con- 
tains all  the  metals  as  sulphids.  Filter  and  wash.  Warm  the  precipitate 
in  yellow  ammonium  sulphid.  Filter  and  wash.  The  nitrate  contains  the 
sulphids  of  As,  Sb,  Sn,  Au,  and  Pt  (Division  B).  The  precipitate  contains 
HgS,  PbS,  Bi2S3,  CuS,  and  CdS  (Division  A). 


METALS  (CATIONS) 


81 


DIVISION  A 

After  the  precipitate,  formed  by  H2S,  has  been  washed,  boil  it  with  HN03 
and  filter. 


A  black^  in- 
soluble resi- 
due   is  ^  HgS, 
often  mixed 
with  sulphur 
and  PbSO4. 
Warm  it  with 
aqua  regia, 
and  after 
filtering     add 
SnCl2.    White 
HgCl   is   pre- 
cipitated, 
changed  to 
gray  Hg  with 
excess   of   the! 
reagent. 


The  solution  contains  Pb,  Bi,  Cu,  and  Cd.  Evaporate  it 
nearly  to  dryness  under  a  hood.  Dilute  with  water  and  pre- 
cipitate the  Pb  with  dilute  H2SO4  and  a  few  drops  of  alcohol, 
if  necessary  for  complete  precipitation.  Filter  and  wash. 


The  precipi- 
tate is  PbSO4. 
Dissolve  in 
NaOH  and  add 
to  the  solution 
HoS.  Black 


PbS      is 
duced. 


pro- 


The  filtrate  contains  Bi,  Cu,  and  Cd. 
Add  NH4OH  in  excess  and,  if  there  is 
a  precipitate,  filter. 


Bi(OH)3  The 

is  the  precipi-  ence  of 


pres- 


m 


Cu 


tate.     After  it  the  solution  is 


on 


has  been 
washed 
moisten  it 
the  paper 
with      NaOH 
then  with 
SnCl2.     It 
turns  brown 
or  black. 


indicated     by  colorless,    then 


the  blue  color. 
Acidify  one 
part  and  add 
K4Fe     (CN).. 
Cu2Fe(CN)6) 
a  reddish- 
brown  com- 
pound 
appears. 


To  the  other 
part  add  KCN 
until  it  is 


H2S.  Yel- 
low CdS  is 
precipitated.  $, 
In  absence 
of  Cu  omit 
KCN. 


DIVISION  B 

Acidify  the  filtrate  slightly  with  HC1.     The  As2S3,  Sb2S8,  SnS2,  Au,S:i, 
and  P-tSr  are  precipitated  with  fine  white  sulphur.     Filter  and  wash,  dis- 
carding the  filtrate.     Warm  the  mixed  sulphids  for 
test-tube  with  (NH4)2CO3.     Filter  and  wash. 


several  minutes  in  a 


The  solu- 
tion contains 
the  As.  Acid- 
ify with  HC1, 
add  H2S;  yel- 
low As2S3 
precipitates. 


The  insoluble  residue  contains  Sb2S3,  SnS2,  Au2S3,  and  PtS2. 
As  far  as  possible  remove  water,  place  in  a  porcelain  dish  and 
warm  und^r  a  hood  with  concentrated  HC1.  Filter. 


The  filtrate  contains  Sb  and 
Sn.  Add  a  little  Zn,  and  touch) 
it  with  a  Pt  foil.  The  Sb 
makes  a  brown  deposit  on  the 
Pt.  The  Sn  is  precipitated  as  a 
heavy  black  powder.  Wash 
this  by  decantation  and  dis- 
solve in  concentrated  HC1.  A! 
few  drops  of  HgCl2  gives  a 
white  precipitate,  HgCll 
(blackened  by  NH4OH),  or 
gray  Hg. 

If  the  foil  is  rinsed,  then! 
warmed  in  (NH4)2Sx,  the  Sb  is 
dissolved  and  after  evapora- 
tion to  dryness  leaves  orange- 
red  Sb2Ss. 


The  insoluble  residue  con- 
tains only  S  and  (possibly) 
As2S3  if  it  is  light-colored,  'if 
Au2S3  or  PtS2  is  present  it  is 
dark-colored.  In  the  latter 
case  warm  it  with  aqua  regia 
and  filter  out  any  insoluble 
residue.  The  filtrate  contains 
Au  and  Pt.  Divide  it  into  two 
portions.  To  one  add  stannous 
and  stannic  chlorids.  A  purple 
color  indicates  gold.  To  the 
other  add  NH4C1.  Yellow  pre- 
cipitate of  (NH4)2PtCl6. 


82  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

Explanation  of  the  Separation  of  the  Cations  of  Group  II 

If  group  I  has  previously  been  removed  from  the  solution 
it  will  contain  hydrochloric  acid,  otherwise  this  must  be 
added,  partly  to  favor  the  precipitation  of  some  of  the  sul- 
phids  in  this  group,  partly  to  prevent  a  precipitation  of  some 
of  those  of  group  III.  A  large  amount  of  hydrogen  sulphid 
will  be  necessary  because  the  solution  cannot  be  made  con- 
centrated and  the  gas  precipitates  slowly. 

Arsenic,  antimony,  and  tin  sulphids  are  all  converted  into 
soluble  sulpho-salts  by  the  ammonium  sulphid.  In  the  case 
of  stannous  sulphid  the  yellow  or  polysulphid  is  necessary. 
This  is  made  by  dissolving  sulphur  in  the  colorless  sulphid, 
when  one  or  more  atoms  of  sulphur  are  added  to  the  mole- 
cule, forming  (NH4)2SX  where  x  =  two  or  more.  The  tin  is 
changed  to  a  stannic  compound  when  it  dissolves. 

SnS-KNH4)2S2=(NH4)2SnS3. 

Acidifying  decomposes  the  sulpho-salt  and  precipitates 
the  tin  as  a  stannic,  not  a  stannous,  sulphid. 


In  the  digestion  of  the  mixed  groups  with  yellow  ammonium 
sulphid  boiling  must  be  prevented  because  this  would  decom- 
pose the  ammonium  sulphid. 

Sulphur  is  always  precipitated,  white  because  of  its  fineness, 
when  a  polysulphid  is  acidified. 

=  2NH4C1+H2S+S. 


Explanation  of  the  Separation  of  the  Cations  of  Group  II, 

Division  A 

The  sulphids  of  copper,  bismuth,  and  cadmium  dissolve 
in  boiling  nitric  acid,  forming  nitrates,  and  lead1  does  so  for 

1  Although  lead  is  classified  with  the  metals  of  group  I,  it  will,  if  present, 
be  found  in  part  in  group  II. 


METALS  (CATIONS)  83 

the  most  part,  but  may  be  partly  oxidized  to  insoluble  lead  sul- 
phate. This  remains  with  the  mercuric  sulphid  which  does 
not  dissolve  in  a  single  acid.  The  lead  compound  does  not 
become  soluble  by  the  action  of  aqua  regia  but  the  mercuric 
sulphid  is  changed  to  the  chlorid.  The  action  of  the  tin 
chlorid  is  as  described  in  87. 

The  portion  of  the  lead  which  went  into  solution  with  the 
other  metals  can  be  precipitated  as  the  sulphate,  but  since  the 
presence  of  much  free  nitric  acid  will  to  some  extent  prevent 
this,  thus  leaving  some  of  the  metal  to  interfere  with  future 
tests,  it  is  better  to  remove  the  acid.  To  accomplish  this  the 
evaporation  must  be  carried  so  far  that  but  a  very  small  bulk 
remains  as  the  acid  is  driven  off  only  after  the  water,  in  the 
last  stages  of  the  operation.  The  lead  sulphate  is  not  en- 
tirely insoluble,  but  can  be  made  more  so  when  alcohol  is 
present.  The  acid  should  be  added  as  long  as  a  precipitate 
appears,  but  a  great  excess  is  objectionable  because  of  the 
large  amount  of  ammonium  hydroxid  which  it  would  neces- 
sitate in  the  next  step. 

A  very  small  quantity  of  copper  gives  the  blue  color  with 
ammonium  hydroxid,  which  must  be  in  excess  to  produce  it. 
If  the  color  is  well-marked  the  ferrocyanid  test  may^be 
unnecessary. 

It  is  necessary  to  confirm  the  bismuth  by  some  other  test 
than  precipitation  by  ammonium  hydroxid  as  if  any  of  the 
lead  had  remained  unprecipitated  by^sulphuric  acid  it  also 
would  form  a  white  precipitate  here.  It  does  not,  however, 
respond  to  the  confirmatory  test  for  bismuth. 

The  addition  of  potassium  cyanid  before  testing  for  cad- 
mium is  to  prevent  the  precipitation  of  the  copper  as  a  sul- 
phid by  hydrogen  sulphid  (69,  82).  This  being  so  much 
darker  than  the  cadmium  sulphid,  the  latter  would  prob- 
ably escape  detection. 


84  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Explanation  of  the   Separation  of  the  Cations  of  Group 
II,  Division  B 

Of  the  precipitated  sulphids  only  that  of  arsenic  is  soluble 
in  ammonium  carbonate,  from  which  solution  it  is  precipi- 
tated as  As2S3  by  acidifying.  Either  Gutzeit's  or  Reinsch's 
test  will  react  with  the  suspended  arsenous  sulphid  without 
its  being  filtered  from  the  liquid. 

The  stannic  sulphid  with  hydrochloric  acid  dissolves  to 
SnCLj,  and  the  antimony  sulphid  to  SbCls.  The  antimony 
compound  is  with  difficulty  attacked  by  the  acid  and  lor  this 
reason  it  must  be  as  far  as  possible  free  from  water  to  avoid 
dilution  of  the  acid.  Boiling  would  drive  off  the  gaseous 
HC1.  The  chlorids  of  tin  and  antimony  are  both  reduced  to 
the  metal  by  nascent  hydrogen.  This  is  produced  by  the 
action  of  an  acid  on  zinc.  If  platinum  is  present  the  anti- 
mony is  deposited  upon  it,  but  the  tin  is  not.  Zinc  almost 
always  contains  impurities  which  remain  after  it  has  dis- 
solved. They  are  often  black  in  color,  but  float  in  the  liquid 
while  the  tin  is  heavy  and  quickly  settles.  To  confirm  it  the 
stannous  chlorid  is  produced  by  the  aid  of  concentrated 
hydrochloric  acid  and  this  is  added  to  2-3  c.c.  of  mercuric 
chlorid,  yielding  white  mercurous  chlorid  or  gray  metallic 
mercury.  The  reactions  are  as  follows: 


2HgCl 

The  fine  coating  of  metallic  antimony  dissolves  readily 
in  yellow  ammonium  sulphid  to  the  sulphantimonite  and  this 
is  decomposed  by  the  heat  of  evaporation  to  antimonous 
sulphid. 

The  sulphids  of  gold  and  platinum  which  do  not  dissolve 
in  a  single  acid  do  so  in  aqua  regia  and  can  be  precipitated 
from  this  by  the  ordinary  reagents. 


METALS  (CATIONS)  85 

Practical  Exercises  in  Group  II 

1.  Mix  in  a  beaker  3-5  c.c.  each  of  solutions  of  mercury, 
lead,  copper,  bismuth,  and  cadmium,  and,  after  precipitating 
with  hydrogen  sulphid,  separate  the  metals  according  to 
Table  IV,  division  A. 

2.  In  the  same  manner  precipitate  a  mixture  of  the  solu- 
tions of  arsenic,  antimony,  and  tin  with  hydrogen  sulphid 
and  separate  the  mixed  sulphids  according  to  Table  IV, 
division  B. 

3.  For  practice  in  separation  of  the  group  into  its  two 
divisions  mix  in  the  same  way  one  or  two  from  each  division, 
e.g.,  mercury,  copper,  arsenic,  and  antimony.     In  making  the 
selection  it  should  be  remembered  that  stannous  chlorid 
reduces  mercuric  compounds  and  that  this  and  other  chlorids 
precipitate  lead  in  part,  also  that  to  keep  bismuth,  and  some- 
times antimony,  in  solution  after  dilution  a  considerable 
quantity  of  acid  is  required. 

4.  For  further  practice  make  analyses  of  unknown  mixtures 
of  the  metals  of  group  II  furnished  by  the  instructors. 

Questions  for  Further  Study  on  Group  II 

What  use  is  made  in  medicine  of  the  fact  that  arsenous  acid 
forms  an  insoluble  compound  with  ferric  or  magnesium 
hydroxid?  How  should  the  ferric  hydroxid  be  prepared  for 
this  purpose  and  what  will  be  the  most  convenient  materials? 
Soluble  salts  of  what  metals  would  be  incompatible  with 
solutions  of  arsenites  and  arsenates  ?  Is  the  insolubility  of  the 
arsenites  and  arsenates  of  the  heavy  metals  sufficient  to  ren- 
der these  compounds  non-poisonous?  What  classes  of  solu- 
ble compounds  would  be  incompatible  with  solutions  of  metals 
of  group  II?  Where  and  how  is  cupric  hydroxid  used  in  test- 
ing for  glucose?  What  is  the  difference  between  the  red  and 


86  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

the  black  mercuric  sulphid?  between  yellow  and  red  mercuric 
oxid?  What  is  the  pharmaceutical  name  of  the  yellow  oxid 
and  how  is  it  prepared?  What  pharmaceutical  names  are 
applied  to  the  mercuric  aminochlorid?  What  medicinal  sub- 
stances have  sufficient  reducing  powers  to  take  the  whole  or  a 
part  of  the  chlorin  from  mercuric  chlorid?  With  what  class 
of  medicinal  compounds  would  a  solution  of  mercuric  chlorid 
in  an  excess  of  potassium  iodid  be  incompatible  ?  What  com- 
pound of  antimony  most  commonly  causes  acute  poisoning? 
Could  any  of  the  above  reagents  be  safely  used  to  render  it 
insoluble  in  the  gastric  juice?  If  so  in  what  form  or  prepara- 
tions could  it  be  most  conveniently  obtained? 

GROUP  I 

Lead,  Silver,  and  Mercury  (Mercurous  Compounds) 

They  are  precipitated  from  their  solutions  by  the  chlorid 
ion.  Their  sulphids,  carbonates,  phosphates,  oxids,  hydrox- 
ids,  arsenates,  arsenites,  bromids,  and  iodids,  as  well  as  their 
chlorids,  are  insoluble  in  water,  although  lead  chlorid  is 

somewhat  soluble. 

Lead,  Pb 

Lead  is  soluble  in  dilute  nitric  acid,  more  easily  by  the  aid 
of  heat,  but  is  little  affected  by  hydrochloric  or  sulphuric 
acids.  Its  soluble  compounds  are  poisonous. 

^A  great  number  of  organic  substances  form  insoluble  com- 
pounds with  lead  acetate,  which  can  consequently  be  used  to 
remove  them  from  solution. 

For  the  reactions  usea  5-per-cent.  solution  of  Pb(C2H3O2)2. 

143.  The  sulphid  ion  gives  with  lead  solutions  a  black  pre- 
cipitate of  lead  sulphid,  PbS,  soluble  in  dilute  nitric  acid,  sul- 
phur being  set  free  at  the  same  time.  If  very  concentrated 
nitric  acid  is  used  sulphuric  acid  is  formed  from  the  sulphur 
and  this  converts  the  lead  into  insoluble,  white,  lead  sulphate, 
PbSO4. 


METALS  (CATIONS)  87 

144.  The  hydroxid  ion  precipitates  lead  hydroxid,  Pb(OH)2- 
This  dissolves  in  an  excess  of  sodium  hydroxid  or  potassium 
hydroxid  with  the  formation  of  a  plumbate  of  the  alkali  metal, 
Na2PbO2  or  K2PbO2.     It  does  not  dissolve  in  an  excess  of 
ammonium  hydroxid. 

145.  The  carbonate  ion  precipitates  white  lead  carbonate. 
Since  the  soluble  carbonates  easily  hydrolyze,  the  reagent 
usually  contains  GET,  and,  consequently,  the  hydroxid  precipi- 
tates with  the  carbonate,  perhaps  as  (PbCO3)2Pb(OH)2.     It 
is  soluble  in  acids,  except  sulphuric. 

146.  The  chlorid  ion  in  concentrated  solutions  of  lead  salts 
forms  'a  heavy  white  precipitate  of  lead  chlorid,  PbCl2.     This 
dissolves  in  hot  water  and  on  cooling,  separates  again  in  the 
form  of  needle-shaped  crystals. 

147.  The  sulphate  ion  gives  a  fine,  heavy,  white  precipitate 
of  lead  sulphate,  PbSO4.     From  very  dilute  solutions  it  only 
appears  after  standing  a  long  time.     Free  acids,  except  sul- 
phuric, hinder  its  separation  and  alcohol  hastens  it.     Lead 
sulphate  dissolves  in  sodium  hydroxid. 

148.  The  chroma  te  or  dichromate  ion  precipitates  yellow 
lead  -chromate,  PbCrO4,  soluble  in  nitric  acid  or  sodium 
hydroxid. 

149.  The  iodra  ion  precipitates  yellow  lead  iodid,  PbI2, 
from  not  too  dilute  solutions.     This  dissolves  in  boiling  water 
from  which  it  separates  as  the  water  cools. 

150.  Tannic  acid  precipitates  yellowish  lead  tannate. 

The  Mercurous  Ion,  Hg* 

Use  for  the  reactions  a  i-per-cent.  solution  of  HgNOs. 

151.  The  sulphid  ion  produces  with  the  mercurous  ion  a 
black  precipitate  of  mercuric  sulphid  mixed  with  mercury. 


,  or 
152.  The  hydroxid  ion  if  obtained  from  compounds  of  so- 


INTRODUCTION   TO   CHEMICAL  ANALYSIS 

dium,  potassium,  or  calcium  forms  a  black  precipitate  of 
mercurous  oxid,  Hg2O,'  insoluble  in  excess. 

153.  Ammonium  hydroxid  precipitates  a  black  mercurous 
amino  salt. 


1  54.  The  chlorid  ion  precipitates  fine,  white,  heavy  mercur- 
ous chlorid,  HgCl.  This  is  soluble  in  aqua  regia,  being 
slowly  converted  into  mercuric  chlorid,  HgCl2.  It  dissolves 
with  much  difficulty  in  nitric  or  hydrochloric  acids  singly. 
Ammonium  hydroxid  changes  it  into  black  mercurous  amino- 
chlorid.  Dry  HgCl  is  slowly  changed  by  sugar  and  some 
other  organic  substances  into  HgCl2  and  Hg. 

155.  The  iodid  ion  precipitates  greenish-yellow  mercurous 
iodid,  Hgl. 

156.  Metallic  copper  and  the  stannous  ion  give  the  same 
reactions  as  with  the  mercuric  compounds.     Other  reducing 
agents  act  like  the  stannous  ion. 

157.  Nitric  acid  and  other  oxidizing  agents  change  mercur- 
ous to  mercuric  compounds. 

Boil  2-3  c.c.  of  mercurous  nitrate  solution  for  a  minute 
with  1-2  c.c.  of  concentrated  nitric  acid.  HC1  will  not  pro- 
duce a  precipitate.  (Compare  154.) 

Silver,  Ag 

Nitric  acid  readily  dissolves  silver,  but  it  does  not  dissolve 
in  hydrochloric  acid.     Its  soluble  salts  are  poisonous. 
Use  for  the  reactions  a  i-per-cent.  solution  of  AgNO3. 

158.  The  sulphid   ion  precipitates    black  silver  sulphid, 
Ag2S,  which  dissolves  in  boiling  nitric  acid. 

159.  The  chlorid  ion  gives  a  precipitate  of  heavy,  curdy, 
white  silver  chlorid,  AgCl.     At  the  time  of  precipitation  it  is 
white,  but  by  exposure  to  the  light  it  gradually  turns  violet, 
then  gray,  and  finally  black.     In  very  dilute  solutions  it  does 
not  precipitate  at  first,  the  liquid  becoming  opalescent  or 


METALS  (CATIONS)  89 

milky.  From  this  the  precipitate  slowly  settles  after  it  has 
stood  for  some  time,  or  more  rapidly  by  boiling.  Silver 
chlorid  is  not  soluble  in  nitric  acid,  but  dissolves  readily  in 
ammonium  hydroxid  and  is  reprecipitated  from  this  solution 
by  acidifying  with  nitric  acid. 

1 60.  The  hydroxid  ion  precipitates  white  silver  hydroxid 
which  quickly  changes  to  the  dark  brown  silver  oxid,  Ag2O. 
This  is  insoluble  in  excess  of  sodium  or  potassium  hydroxid 
but  with   ammonium   hydroxid  it  forms   a  complex  ion, 
Ag(NH2)3',  becoming  soluble. 

161.  The  chromate  or  dichromate  ion  forms  a  dark  red 
precipitate  of  silver  chromate,  Ag2CrO4. 

162.  Zinc,  copper,  mercury,  tin,  or  lead  precipitate  the  metal. 

163.  Ferrous  sulphate,  FeSO4,  and  other  reducing  agents 
convert  the  silver  ion  into  the  metal. 

Separation  of  the  Cations  of  Group  I 

Add  hydrochloric  acid  to  the  solution.  Silver  and  mercury 
(ous)  are  precipitated  with  most  of  the  lead  as  chlorids. 
Filter  and  wash  in  cold  water.  Make  a  hole  in  the  apex  of 
the  filter  and  rinse  the  precipitate  with  10-15  c-c-  °f  water 
into  a  test-tube.^  Heat  to  boiling  and  filter  while  hot.  The 
lead  chlorid  is  cffl^Bed.  Test  a  portion  of  the  nitrate  after  it 
is  cold  for  this  with  dilute  sulphuric  acid.  A  fine  white  pre- 
cipitate of  lead  sulphate  is  formed.  From  another  portion 
with  potassium  dichromate  precipitate  yellow  lead  chromate. 

After  the  precipitated  chlorids  of  silver  and  mercury  have 
been  washed  with  hot  water  on  the  filter,  pour  over  it  10  c.c. 
of  warm  ammonium  hydroxid,  receiving  the  filtrate  in  a  clean 
test-tube.  The  silver  chlorid  dissolves.  Mercurous  chlorid 
is  changed  to  a  black  solid  which  remains  upon  the  filter. 

Acidify  the  filtrate  with  nitric  acid.  The  silver  chlorid 
appears  as  a  white  precipitate  or  opalescence,-  soluble  in 
ammonia  and  reprecipitated  by  nitric  acid. 


QO  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

TABLE  V 

OUTLINE  OF  SEPARATION  or  THE  CATIONS  OF  GROUP  I 

To  the  solution  of  the  mixed  ions  add  HC1  in  excess.  Filter  and  wash 
in  cold  water.  The  precipitate  contains  PbCl2,  AgCl,  and  HgCl.  All 
other  groups  present  in  the  original  solution  pass  into  the  nitrate.  Boil  the 
precipitate  with  water,  filtering  hot. 


The  nitrate  contains  Pb. 
Test  part  with  H2SO4;  a  fine 
white  precipitate  is  lead  sul- 
phate. 

Test  part  with  K2Cr2O7;  a 
yellow  precipitate  is  lead  chro- 


mate. 


The  insoluble  residue  contains  AgCl  and  HgCl. 
Dissolve  the  AgCl  in  warm  NH4OH. 


Acidify  the  filtrate  with 
nitric  acid;  a  white  pre- 
cipitate is  AgCl. 


The  mercury  re- 
mains as  black,  in- 
soluble NHjHgaCl. 


Explanation  of  the  Separation  of  the  Cations  of  Group  I 

This  depends  upon  the  solubility  of  the  chlorids,  lead 
chlorid  dissolving  in  hot  water  but  not  easily  in  cold,  and 
silver  chlorid  in  ammonium  hydroxid  and  not  in  hot  water, 
the  mercurous  chlorid  being  insoluble  in  both. 

The  lead  can  be  confirmed  in  the  aqueous  solution  by  any 
of  its  characteristic  precipitants. 

With  ammonium  hydroxid  silver  chlorid  forms  ammonio- 
silver  chlorid,  Ag(NH3)2Cl,  which  is  soluble.  It  is  decom- 
posed by  acids  the  silver  chlorid  separating  again. 


The  identification  of  the  mercurous  chlorid  is  without  diffi- 
culty since  it  is  the  only  one  of  the  three  which  gives  a  dark 
color  with  ammonium  hydroxid. 

Practical  Exercise  on  Group  I 

Mix  in  a  test-tube  4-5  c.c.  each  of  solutions  of  silver,  lead, 
and  mercurous  ions  and  make  the  separation  as  given  above. 

Questions  for  Further  Study  on  Group  I 

Wouft  trie  compounds  of  lead  which  are  formed  as  pre- 
cipitates be  necessarily  insoluble  in  the  gastric  juice?  Are 


METALS  (CATIONS)  91 

there  any  of  the  reagents  which  might  be  employed  in  cases  of 
acute  lead  poisoning  as  antidotes?  Would  all  of  those  which 
form  insoluble  compounds  and  which  would  thus  prevent  ab- 
sorption, be  suitable  for  such  a  purpose?  What  organic  com- 
pounds commonly  used  medicinally  would  be  incompatible 
with  lead  compounds?  What  is  the  common  name  of  lead 
carbonate?  In  accordance  with  the  above  reactions  what 
would  be  a  safe  and  easily  obtainable  antidote  for  a  poisonous 
dose  of  mercurous  or  silver  nitrate?  What  is  the  pharma- 
ceutical name  of  mercurous  chlorid?  How  does  it  compare 
in  its  chemical  properties  and  physiological  action  with 
mercuric  chlorid?  Would  the  chemical  antidotes  for  the 
mercurous  ion  necessarily  precipitate  mercuric  chlorid? 
What  is  the  danger  of  leaving  dry  mercurous  chlorid  for  a 
long  time  mixed  with  organic  substances  ?  What  is  a  simple 
method  of  testing  calomel  for  small  quantities  of  mercuric 
chlorid?  What  is  the  pharmaceutical  name  of  mercurous 
oxid  and  how  is  it  prepared?  Where  is  the  turning  dark  of 
silver  chlorid  by  light  made  use  of  practically? 

The  Separation  of  Metals  into  Groups 

The  different  metals  can  be  precipitated  in  succession  in 
groups  from  the  original  solution  by  using  these  group  re- 
agents in  the  following  order: 

1.  Hydrochloric  acid,  which  precipitates  group  I,  silver, 
the  mercurous  ion  and  lead  (in  part) . 

2.  Hydrogen  sulphid,  which  precipitates  group  II,  lead 
(remainder),  the  mercuric  ion,  copper,  bismuth,  cadmium, 
arsenic,  antimony,  tin,  gold,  and  platinum. 

3.  Ammonium   sulphid    (in   the   absence  of  phosphates, 
oxalates,  borates,  citrates,  and  tartrates),  which  precipitates 
group  III,  cobalt,  nickel,  iron,  manganese,  chromium,  zinc, 
and  aluminum. 


Q2  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

4.  Ammonium  carbonate,   which  precipitates  group  IV, 
barium,  strontium,  and  calcium,  but  not  magnesium. 

Magnesium,  sodium,  potassium,  and  ammonium  are  not 
precipitated  by  any  of  the  above  reagents. 

fine  details  of  the  separation  are  as  follows: 

|To  the  cold  solution  which  may  contain  all  the  ions  add  hy- 
drochloric acid1  as  long  as  any  precipitate  forms.  This  con- 
tains group  I.  Separate  the  metals  according  to  Table  V. 
Filter,  and  to  the  nitrate  add  hydrogen  sulphid  until  the  liquid 
smells  of  this  gas.  Warm,  stir  to  make  the  precipitate  collect 
in  large  masses,  then,  if  the  precipitation  is  complete,  filter  and 
wash  with  warm  water.  The  metals  of  group  II  are  on  the 
filter,  the  remainder  of  the  groups  in  the  filtrate.  Make  the 
separation  according  to  Table  IV.  If  the  aqueous  solution  of 
hydrogen  sulphid  has  been  used  to  precipitate  group  II,  and 
the  filtrate  is  therefore  of  large  volume,  it  may  be  concentrated 
to  40-50  c.c.  by  boiling.  It  should  then  be  made  neutral  or 
faintly  alkaline  by  ammonium  hydroxid  and,  after  the  addi- 
tion of  ammonium  chlorid,  be  precipitated  by  a  slight  excess  of 
ammonium  sulphid.  Warm  nearly  to  boiling,  filter,  and  wash. 
The  precipitate  contains  the  metals  of  group  III.  Separate 
them  as  in  Table  II  or  III.  Boil  the  filtrate  until  the  ammo- 
nium sulphid  has  been  decomposed,  as  is  shown  by  the  dis- 
appearance of  its  odor,  filtering  out  the  sulphur  if  any  has 
been  set  free.  Precipitate  group  IV  with  ammonium  carbon- 
ate, warming  nearly  to  boiling  to  insure  its  complete  separa- 
tion. Filter  and  wash.  Analyze  the  group  by  the  use  of 
Table  I.  If  the  filtrate  is  much  diluted,  evaporate  to  10  or  20 
c.c.  and  apply  the  flame  tests  for  sodium  and  potassium. 
Afterward  add  to  the  solution  sodium  phosphate  as  a  test 
for  magnesium.  Test  the  original  substance  for  ammonium. 

1  In  precipitating  the  different  groups,  when  their  presence  or  absence  is 
uncertain,  it  is  usually  better  to  first  test  only  a  small  portion  of  the  solution 
with  the  group  reagent,  then,  if  a  precipitate  forms,  to  treat  the  whole  in  the 
same  manner. 


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94  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

The  Method  of  Analysis  of  Alloys  Containing 

Hg,  Ag,  Cu,  Bi,  Pb,  Cd,  Au,  Pt,  Sb,  Sir,  Ni,  Co,  Fe, 
Mn,  Cr,  Al,  Zn 

Warm  the  fine  alloy  with  equal  parts  of  concentrated  nitric 
acid  and  water  under  a  hood  until  no  gritty  particles  re- 
remain.  Au,  Pt,  Sb  and  Sn  are  not  dissolved;  the  others  go 
into  solution.  Filter  and  wash. 

Evaporate  the  filtrate  nearly  to  dryness  if  much  acid  has 
been  used,  dilute,  heat  to  boiling  and  precipitate  Ag  while  hot, 
by  HC1,  filter  and  precipitate  by  an  excess  of  H2S,  which 
forms  insoluble  sulphids  of  Hg,  Cu,  Bi,  Pb  and  Cd.  Stir 
until  it  settles,  then  filter  and  separate  metals  in  the  filtrate 
by  Table  II.  Wash  the  precipitate  well,  then  rinse  it  from 
the  paper  with  about  10  c.c.  of  nitric  acid  and  separate  the 
metals  by  Table  IV,  Division  A. 

If  Au  or  Pt  are  contained  in  the  first  insoluble  residue  it 
will  be  dark  in  color.  The  compounds  of  Sb  and  Sn  are 
white.  Dissolve  the  moist  precipitate  in  a  little  warm  aqua 
regia  and,  after  filtering,  boil  the  solution  a  minute  under  the 
hood  to  expel  the  chlorin.  Test  half  for  Sn  and  Sb  by  metallic 
zinc  and  platinum  (116,  127)  (Table  IV,  Division  B).  Evap- 
orate the  excess  of  acid  from  the  other  half  of  the  solution 
under  a  hood  and  dissolve  in  3-4  c.c.  of  water.  Test  a  part 
for  Au  by  SnCl2  and  SnCl4,  which  gives  a  purple  color 
(purple  of  Cassius).  To  the  rest  add  concentrated  ammo- 
nium chlorid.  Pt  gives  a  yellow  crystalline  precipitate  (140) . 

Practical  Exercises  on  General  Analysis 

For  further  drill  the  student  should  make  analyses  of  as 
many  unknown  substances  as  possible,  where  all  the  groups 
must  be  sought  for. 


METALS  (CATIONS)  95 

Questions  for  Further  Study  on  the  Metallic  Compounds 

Which  of  the  metals  are  incompatible  with  the  alkaline 
hydroxids  because  of  their  being  precipitated  ?  Does  calcium 
hydroxid  have  a  similar  action  to  that  of  the  hydroxids  of  the 
alkali  metals?  What  hydroxids  have  an  alkaline  reaction? 
What  carbonates?  What  ones  of  the  common  pharmaceu- 
tical preparations  have  an  alkaline  reaction  ?  Which  groups 
of  metals  are  precipitated  by  soluble  carbonates?  Of  which 
ones  are  the  carbonates  soluble  in  water?  What  medicinal 
substances  with  an  oxidizing  action  are  the  most  commonly 
used  in  solution?  What  ones  have  a  reducing  action? 
What  compounds  of  the  metals  studied  would  be  so  far  af- 
fected by  the  last-mentioned  reagents  as  to  cause  an  incom- 
patibility? What  classes  of  compounds  would  be  incom- 
patible with  the  metallic  compounds  of  group  I  by  rendering 
them  insoluble?  with  group  II?  with  group  III?  with 
group  IV?  with  group  V?  What  pharmaceutical  prepara- 
tions contain  tannic  acid  ?  Does  albumin  form  insoluble  com- 
binations with  any  other  metallic  compounds  except  mercuric 
salts?  ,Are"any  compounds  decomposed  by  exposure  to  light 
except  silver  chlorid?  What  is  the  usual  composition  of  pre- 
cipitates produced  by  the  action  of  a  soluble  hydroxid  on  the 
ion  of  a  soluble  metallic  salt?  by  a  soluble  carbonate?  What 
are  the  exceptions  to  the  usual  rule?  Knowing  the  composi- 
tion of  the  reagent  used,  sulphid,  phosphate,  chromate,  sul- 
phate, etc.,  can  we  predict  what  will  be  the  composition  of  the 
precipitate  which  will  be  formed  with  the  different  cations? 


96  INTRODUCTION   TO   CHEMICAL  ANALYSIS 


CHAPTER  II 

ACIDS  (ANIONS) 

IN  testing  for  the  presence  of  an  acid  we  must  distinguish 
between  the  acid  in  the  free  state  and  that  which  is  combined 
as  a  salt  (the  hydrogen  being  wholly  or  in  part  replaced  by  a 
metal) ,  as  the  reactions  of  one  form  often  differ  from  those  of 
the  other.  In  the  second  case  the  reactions  are  due  to  the 
anion ;  in  the  first  there  are,  in  addition,  those  of  hydrogen  ion. 
Free  acids  usually  turn  blue  litmus  red,  but  the  same  is  true  of 
many  acid  salts  as  well  as  of  some  normal  salts  of  the  heavy 
metals  (pp.  20,  21).  The  acids  may  be  classified  in  a  number 
of  groups  according  to  the  similarity  of  their  chemical  action, 
but  the  actual  separation  of  the  anions  from  a  complex  mix- 
ture is  not  so  simple  a  proceeding  as  that  with  the  cations. 

Sulphuric  Acid,  H2SO4,  and  Sulphates  (the  Sulphate  Ion) 

164.  Pure  sulphuric  acid  is  a  heavy,  oily  liquid  (sp.  gr.  1.85) 
which  generates  a  large  amount  of  heat  when  diluted  with 
water,  with  which  it  forms  a  number  of  compounds.  It 
dissolves  many  metals  as  sulphates. 

The  free  acid  chars  most  organic  compounds  when  it  is 
warmed  with  them  by  abstracting  the  elements  of  water 
and  leaving  the  carbon.  If  the  acid  is  dilute,  the  charring 
only  occurs  after  the  water  has  been  evaporated,  leaving  the 
concentrated  acid.  This  can  be  shown  upon  paper  or  wood. 
Some  organic  compounds,  including  oxalic  acid  and  its  salts, 
also  some  of  the  alkaloids,  are  not  thus  blackened. 

For  the  reactions  of  the  anion  a  i-per  cent,  solution  of  any 
sulphate  may  be  used. 


ACIDS  (ANIONS)  97 

165.  The  barium  ion  gives,  in  solutions  of  the  free  acid  or 
its  salts,  a  white,  heavy,  finely  divided  precipitate  of  barium 
sulphate,   BaSO4,  which  is  insoluble  in  water,    acids^and 
alkalies. 

1 66.  The  lead  ion  precipitates  the  sulphate  ion  as  heavy, 
white,  finely  divided  lead  sulphate,  PbSO4.     It  is  somewhat 
soluble  in  acids  and,  after  washing,  blackens  when  brought 
in  contact  with  ammonium  sulphid.     It  is  not  so  sensitive  a 
test  as  the  preceding. 

Sulphurous  Acid,  H2SO3,  and  Sulphites  (the  Sulphite  Ion) 

Sulphurous  acid  is  produced  when  sulphurous  oxid  is  dis- 
solved in  water.  It  is  a  colorless  liquid  with  the  odor  of  burn- 
ing sulphur.  It  has  an  acid  reaction  and  bleaches  unstable 
coloring  matters  like  those  of  the  flowers.  By  exposure  to 
the  air  it  is  oxidized  readily  to  sulphuric  acid.  Its  salts 
undergo  the  same  chemical  change.  Both  are  consequently 
used  as  reducing  agents.  Only  the  sulphites  of  the  alkalies 
are  soluble  in  water. 

For  the  reactions  use  sodium  sulphite. 

167.  Sulphites  are  decomposed  by  sulphuric  acid  with  the 
evolution  of  SO2,  a  heavy  colorless  gas  which  smells  like  burn- 
ing sulphur.     If  a  large  test-tube  is  allowed  to  fill  with  it, 
after  adding  a  little  acid  to  a  crystal  of  the  sulphite,  and  the 
gas  is  then  poured  into  another  tube  containing  dilute  potas- 
sium permanganate  solution,  the  latter  is  decolorized. 

Oxidizing  agents  like  nitric  acid,  bromin  or  chlorin  water, 
and  many  others  convert  sulphites  into  sulphates,  which  then 
give  with  the  barium  ion  a  precipitate  that  is  insoluble  in  acids. 

1 68.  The  barium  ion  gives  with  a  neutral  solution  of  a  sul- 
phite a  white  precipitate  of  barium  sulphite,  BaSOs,  which 
dissolves  in  hydrochloric  or  nitric  acids. 

169.  The  silver  ion  precipitates  the  sulphite  ion  as  white 
silver  sulphite,  AgSOs,  which  upon  warming  turns  black  with 


98  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

a  formation  of  metallic  silver  and  sulphur  dioxid  and  silver 
sulphate. 

Ag2+SO2+Ag2SO4,  or 


170.  The   sulphite   ion  evolves  hydrogen  sulphid  when 
warmed  with  dilute  sulphuric  acid  and  zinc.     This  gas  turns 
lead  acetate  paper  brown  or  black,  and  can  be  identified  by  its 
odor  likewise. 

Carbonic  Acid,  H2CO3,  and  Carbonates  (the  Carbonate  Ion) 

Carbon  dioxid  (carbon  anhydrid)  when  dissolved  in  water 
forms  carbonic  acid,  H2CO3,  which  reddens  litmus-paper. 
The  acid  exists  only  in  salts  and  in  solution,  which  latter  is 
readily  decomposed  by  heating,  the  carbon  dioxid  being  ex- 
pelled. Of  the  carbonates  only  those  of  the  alkalies  are  very 
soluble  in  water  and  these  when  dissolved  have  in  conse- 
quence of  hydrolysis  a  strong  alkaline  reaction. 

For  the  reactions  sodium  carbonate  may  be  used. 

171.  Carbonates  are  decomposed  by  almost  all  free  acids, 
usually  without  heating,  although  this  is  sometimes  neces- 
sary.    The  gas  which  escapes  with  effervescence  is  carbon 
dioxid.     It  can  be  identified  by  its  rendering  lime-water 
milky.     It  is  a  colorless,  odorless  gas,  so  much  heavier  than 
air  that  it  can  be  poured  from  the  test-tube  where  it  is  evolved 
into  another  containing  2  or  3  c.c.  of  lime-water.     If  the  tube 
is  now  shaken  so  as  to  mix  the  two,  the  liquid  is  turned  white 
—  a  characteristic  reaction.     It  may  also  be  shown  by  hold- 
ing a  drop  of  the  lime-water  down  in  the  tube  of  the  gas  on 
the  end  of  a  glass  rod,  being  careful  not  to  let  the  drop  run 
off  by  touching  the  side  of  the  tube.     The  precipitate  dis- 
solves in  HC1. 

172.  The  acid  carbonates  or  bicarbonates,  like  NaHCO3, 
are  less  stable  than  the  normal  carbonates,  like  Na2CO3  or 


ACIDS  (ANIONS)  99 

CaCOs,  and  are  decomposed  when  heated  in  water  before 
this  reaches  its  boiling-point,  half  the  carbon  dioxid  escaping. 
The  normal  carbonates  are  not  thus  decomposed.  Soluble 
normal  carbonates  will  also  turn  yellow  turmeric  paper  red- 
dish-brown, while  the  acid  carbonates  do  not  if  they  contain 
none  of  the  normal  compound,  although  this  is  very  com- 
monly present. 

173.  The  barium  ion,  with  carbonates  soluble  in  water, 
gives  a  white  precipitate  of  barium  carbonate,  BaCOs.     In 
very  dilute  solutions  it  appears  only  after  heating  to  boiling. 
It  is  soluble  in  most  acids  except  sulphuric.     No  precipitate 
is  produced  with  the  free  carbonic  acid. 

Oxalic  Acid,  H2C2O4,  and  Oxalates  (the  Oxalate  Ion) 

The  acid  is  a  white  solid,  readily  soluble  in  water  and 
crystallizing  from  its  solution  with  two  molecules  of  water 
of  crystallization.  It  has  an  intensely  sour  taste  and  is 
strongly  acid  to  litmus. 

174.  Pure  oxalic  acid  volatilizes  without  leaving  a  residue 
when  heated  on  a  platinum  foil,  giving  irritating  fumes. 
Oxalates  of  the  non-volatile  bases  treated  in  the  same  manner 
are  partly  vaporized,  but  leave  a  residue  behind,  the  carbonate 
of  the  base.   ^>  \\ 

175.  When  heated  nearly  to  boiling  with  concentrated 
sulphuric  acid  the  solid  oxalic  acid  or  an  oxalate  is  decom- 
posed yielding  carbon  dioxid  and  carbon  monoxicj.     The 
former  can  be  identified  by  the  turbidity  which  it  produces 
in  lime-water  and  the  latter  by  the  bright  blue  flame  when 
it  is  ignited.     Oxalic  acid  is  not  blackened  by  concentrated 
sulphuric  acid. 

For  reactions  of  the  oxalate  ion  ammonium  oxalate  may  be 
employed. 

176.  The  barium  ion  precipitates  from  neutral  solutions 


100  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

of  the  oxalates  white,  pulverulent,  barium  oxalate,  BaC2O4. 
It  is  insoluble  in  acetic  acid.  It  dissolves  in  hydrochloric  acid 
and  is  reprecipitated  from  this  by  ammonia.  The  calcium  ion 
gives  a  similar  result,  even  in  very  dilute  solutions. 

177.  Potassium  permanganate  solution  is  quickly  decolor- 
ized when  warmed  with  oxalic  acid  or  one  of  its  salts,  after 
acidifying  with  sulphuric  acid.  This  is  because  of  the  strong 
reducing  power  of  the  acid.  For  the  same  reason  an  acidified 
solution  of  potassium  dichromate  is  likewise  reduced  to  a  green 
chromic  salt. 


Chromic  Acid,  H2CrO4,  and  Chromates  (the  Chromate  Ion) 

Chromic  anhydrid,  CrOs,  when  dissolved  in  water  forms 
chromic  acid,  which  only  exists  in  solution.  The  acid  and  the 
chromates  are  all  colored  red  or  yellow.  The  normal  chro- 
mates  like  K2CrO4  which  contain  two  univalent  metallic  atoms 
to  each  acid  radical,  are  yellow,  but  upon  acidifying  they  be- 
come reddish  by  the  abstraction  of  a  part  of  the  base  and  the 
formation  of  bichromates,  also  called  dichromates,  like 
K2Cr207.  Chromic  anhydrid  and  chromates  act  as  oxidizing 
agents,  the  chromium  taking  a  positive  charge  and  becoming 
the  positive  element  of  the  product,  if  free  mineral  acids  are 
present. 

For  the  reactions  use  a  2-per-cent.  solution  of  potassium 
dichromate. 

178.  The  barium  ion  precipitates  from  solutions  of  chro- 
mates yellow  barium  chromate,  BaCrO-i,  insoluble  in  acetic 
but  soluble  in  hydrochloric  acid. 

179.  The  lead  ion  precipitates  lemon-yellow  lead   chro- 
mate, PbCrO4,  soluble  with  difficulty  in  nitric  acid,  readily 
soluble  in  sodium  hydroxid, 

1 80.  The  sulphid  ion  slowly  changes  an  acidified  solution  of 
a  chromate  to  brown,  then  green,  chromic  salt,  without  pre- 


ACIDS  (ANIONS)  101 

cipitation.     If  the  solution  is  alkaline  there  is  precipitated 
upon   warming   bluish-green   chromic   hydroxid,    Cr(OH)s. 

181.  Concentrated  hydrochloric  acid  evolves  chlorin  when 
warmed  with  a  chromate,  the  chromium  changing  from  the 
anion  to  the  cation.     The  color  of  the  solution  is  changed 
from  yellow  to  green. 

182.  Alcohol  or  oxalic  acid  is  oxidized  by  a  solution  of  potas- 
sium dichromate  acidified  strongly  with  sulphuric  acid,  the 
chromium  changing  as  in  the  preceding  test  and,  conse- 
quently, the  color  of  the  solution  being  changed  to  green. 

Boric  Acid  (Boracic  Acid),  HgBOa,  and  Borates  (the  Borate  Ion) 

The  free  acid  is  a  white,  crystalline  solid,  soluble  in  water. 
It  reddens  litmus-paper  and  turns  to  a  reddish-brown  paper 
which  has  been  colored  yellow  by  a  solution  of  turmeric.  It 
is  somewhat  volatile  with  boiling  water  or  alcohol  and  when 
the  latter  is  ignited  it  imparts  a  green  border  to  the  flame. 
The  borates  of  metals  except  those  of  the  alkalies  are  insoluble 
in  water,  but  soluble  in  most  acids. 

For  the  reactions  a  2-per-cent.  solution  of  borax  may  be  used. 

1 83 .  The  barium  ion  precipitates,  except  in  dilute  solutions, 
white  barium  borate. 

184.  A  few  grains  of  the  solid  moistened  with  a  few  drops 
of  concentrated  sulphuric  acid  in  a  porcelain  dish  will  give  a 
green-bordered  flame  when  mixed  with  2-3  c.c.  of  alcohol  and 
the  latter  is  set  on  fire,  and  stirred  with  a  glass  rod  while  burn- 
ing.    The  free  acid  does  the  same  without  the  addition  of 
sulphuric  acid  (compare  with  75). 

185.  Slightly  acidify  the  solution  of  a  borate  with  hydro- 
chloric acid,  dip  into  this  a  piece  of  turmeric  paper  and  dry 
the  latter  at  a  gentle  heat.     It  is  turned  reddish-brown. 
Moistening  this  with  an  alkali  turns  it  bluish  or  greenish-black. 

1 86.  From  a  hot,  concentrated  solution  of  a  borate  hydro- 


102  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

chloric  acid  precipitates  free  boric  acid  which  separates  in 
white  crystalline  scales. 


Orthophosphoric  Acid,  H-jPC^,  and  Orthophosphates  (the 
Orthophosphate  Ion) 

This  is  the  most  common  acid  of  phosphorus  and  is  the 
one  which  is  meant  when  phosphoric  acid  is  spoken  of  without 
any  distinguishing  prefix.  Although  it  can  be  obtained  in 
the  crystalline  form,  it  is  very  deliquescent,  so  that  it  usually 
appears  as  a  liquid,  thick  and  syrupy  when  concentrated.  It 
has  a  strong  acid  reaction  and  taste. 

For  testing  use  a  2-per-cent.  solution  of  sodium  phosphate, 
Na2HPO4. 

187.  The  barium  ion  precipitates  from  solutions  of  phos- 
phates white  barium  phosphate,  BaHPO4,  soluble  in  most 
acids  except  sulphuric.  It  does  not  produce  this  with  solu- 
tions of  the  free  acid. 

1  88.  Magnesia  mixture  (the  clear  solution  produced  by 
adding  to  magnesium  sulphate,  ammonium  chlorid,  then 
making  it  alkaline  with  ammonium  hydroxid)  precipitates 
from  solutions  of  phosphates  white  ammonium  magnesium 
phosphate,  NH4MgPO4.  Under  the  microscope  it  is  seen 
to  be  in  the  form  of  snowflake-shaped  crystals.  The  com- 
pound is  soluble  in  acids.  In  dilute  solutions  the  precipita- 
tion is  slow  and  is  favored  by  shaking  or  stirring. 

189.  If  0.5  c.c.  of  a  solution  of  the  phosphate  ion  is  added 
to  5  c.c.  of  ammonium  molybdate  in  nitric  acid  and  the  mix- 
ture slightly  warmed,  a  yellow  precipitate  of  ammonium 
phosphomolybdate,  (NH4)3PO4(MoO3)io,  2  H2O  forms  slowly. 
It  is  easily  soluble  in  ammonium  hydroxid,  and  its  formation* 
is  hindered  by  the  presence  of  chlorids  and  of  some  organic 
compounds. 

190.  From  solutions  of  phosphates  containing  no  free  acid, 
the  ferric  ion  precipitates  yellowish-white,  gelatinous  ferric 


ACIDS  (ANIONS)  103 

phosphate.  If  a  little  mineral  acid  is  present,  it  can  be  re- 
moved by  the  addition  of  a  few  crystals  of  sodium  acetate. 
An  excess  of  the  ferric  ion  should  then  be  avoided  as  it  pro- 
duces reddish  ferric  acetate. 

191.  A  solution  of  egg- albumin  is  not  precipitated  by  or- 
thophosphoric   acid   nor   by   solutions   of   orthophosphates 
acidified  with  acetic  acid. 

Metaphosphoric  Acid,  HPO3,  and  Metaphosphates  (the 
Metaphosphate  Ion) 

The  free  acid  is  a  transparent,  ice-like  solid,  readily  soluble 
in  water.  Boiling  in  water  changes  it  into  orthophosphoric 
acid.  The  solution  when  mixed  with  zinc  oxid  yields  a 
plastic  mass  which  hardens  on  standing. 

A  2-per-cent.  solution  of  sodium  metaphosphate  can  be 
used  for  testing. 

192.  The  barium  ion  precipitates  from  neutral  solutions 
barium  metaphosphate,  Ba(POs)2,  soluble  in  hydrochloric 
acid. 

193.  Magnesia  mixture  (188)  with  the  metaphosphate  ion 
gives  no  precipitate  that  is  insoluble  in  ammonium  chlorid. 

194.  The  molybdate  ion  also  fails  to  form  a  precipitate. 

195.  Free  metaphosphoric  acid,  or  solutions  of  metaphos- 
phates  when  acidified  with  acetic  acid,  will  precipitate  a  solu- 
tion of  egg-albumin.     (Difference  from  orthophosphoric  acid.) 

Hypophosphorous  Acid,  HPH2O2,  and  Hypophosphites  (the 
Hypophosphite  Ion) 

The  hypophosphites  are  nearly  all  soluble  in  water.  By 
heating  the  dry  salts  they  are  decomposed  into  phosphates 
and  hydrogen  phosphid,  PH3,  a  combustible  gas  which  be- 
fore burning  has  the  odor  of  decaying  fish. 

2NaPH202  =  PH3  +  Na2HPO4. 
Oxidizing  agents  convert  them  into  phosphates. 


IO4  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Use  for  the  reactions  a  2-per-cent.  solution  of  NaPH2O2. 

196.  The  silver  ion  precipitates,  first,  white  silver  hypo- 
phosphite,  AgPH2O2,  from  which  black  metallic  silver  soon 
separates,  the  latter  change  being  hastened  by  warming. 

2NaPH202+  2AgN03  =  2  Ag+  2NaH2PO4+  N2O2,  or 

2Na'  +  2PH2(Y+2Ag  +  2N08'  =  2Ag+2Na'  +  2H2PO4' 
+  N202. 

197.  The  mercuric  ion  from  mercuric  chlorid  slowly  forms 
white  insoluble  mercurous  chlorid,  HgCl,  more  rapidly  on 
heating. 


Long  heating  changes  the  color  of  the  chlorid  to  gray  from 
the  separation  of  metallic  mercury. 

198.  Hypophosphites    readily    decolorize    a    solution    of 
KMnO4. 

199.  When  warmed  with  concentrated  sulphuric  acid  the 
hypophosphites  cause  an  evolution  of  sulphurous  oxid. 

200.  The  barium  and  lead  ions  do  not  precipitate  the 
hypophosphite  ion. 

All  these  reactions  show  the  reducing  power  of  the  hypo- 
phosphite  ion. 

Thiosulphuric  Acid,  H2S2O3,  and  Thiosulphates  ("Hyposulphites") 
(the  Thiosulphate  Ion) 

The  acid  does  not  exist  free.     Its  salts  are  mostly  soluble 
and  act  as  reducing  agents. 

Sodium  thiosulphate  can  be  used  for  the  reactions. 

201.  The  silver  ion  produces  a  white  precipitate  of  silver 
thiosulphate,  Ag2S2O3.     It  dissolves  in  an  excess  of  the  thio- 
sulphate and  is  therefore  not  readily  precipitated  from  con- 
centrated solutions  of  the  latter  salt.     When  heated  it  turns 
black,  silver  sulphid  being  precipitated. 


ACIDS  (ANIONS)  105 

202.  The    barium    ion   precipitates   white   barium   thio- 
sulphate,  BaS2O3,  which  is  somewhat  soluble  in  water. 

203.  Solutions  of  thiosulphates  are  decomposed  by  the 
mineral  acids  with  an  evolution  of  sulphurous  oxid,  sulphur 
being  at  the  same  time  precipitated. 

204.  An  acidified  solution  of  potassium  permanganate  is 
immediately  decolorized  by  a  thiosulphate. 

Hydrosulphuric  Acid  (Hydrogen  Sulphide),  H2S,  and  Sulphids 
.    (the  Sulphid  Ion) 

Hydrosulphuric  acid  is  a  colorless  gas  with  the  odor  of  de- 
caying eggs.  It  dissolves  in  about  one-third  of  its  volume  of 
water  at  ordinary  temperatures  and  this  solution  reddens 
litmus-paper.  By  the  addition  of  alkaline  hydroxids  to  this 
sulphids  are  formed. 

2NH4OH+H2S  =  (NH4)2S+2H20. 

Of  its  salts  only  those  of  the  alkalies  and  alkaline  earths  are 
soluble  in  water.  These,  together  with  the  sulphids  of  iron, 
manganese,  and  zinc  evolve  hydrogen  sulphid  when  treated 
with  sulphuric  acid.  The  sulphids  of  some  of  the  heavy 
metals  like  copper,  mercury,  and  lead  give  sulphurous  oxid  in- 
stead. Solutions  of  the  sulphids  dissolve  sulphur  with  the 
formation  of  the  polysulphids  — 

(NH4)2S+S  =  (NH4)2S2,  etc. 

For  the  reactions  a  i-per-cent.  solution  of  Na2S  or  (NH4)2S 
can  be  used. 

205.  The  lead  ion  or  silver  ion  with  the  sulphid  ion  form 
black  precipitates  of  lead  sulphid,  PbS,  or  silver  sulphid,  Ag2S. 

206.  Soluble  sulphids  when  acidified  with  hydrochloric  or 
sulphuric  acid  give  hydrogen  sulphid. 


This  can  be  identified  by  its  odor  or  by  its  turning  brown  or 
black  a  paper  previously  dipped  in  a  solution  of  lead  acetate. 


IO6  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

207.  Polysulphids  give  the  same  gas  when  acidified,  and 
also  precipitate  sulphur,  very  fine  and  nearly  white. 

208.  Mercuric  sulphid,  lead  sulphid,  and  some  of  the  other 
insoluble  sulphids  yield  no  hydrogen  sulphid  when  warmed 
with  hydrochloric  acid,  but  do  so  if  at  the  same  time  a  frag- 
ment of  zinc  is  present. 

Hydroferrocyanic  Acid,  H4Fe(CN)6,  and  Ferrocyanids 
(the  Ferrocyanid  Ion) 

Use  for  the  reaction  a  i-per-cent.  solution  of  K4Fe(CN)6. 

209.  The  ferric  ion  forms  a  deep  blue  precipitate  of  ferric 
ferrocyanid  (55). 

210.  The  copper  ion  produces  a  reddish-brown  precipitate 
of  copper  ferrocyanid  (73). 

211.  Ferrocyanids    heated    with    concentrated    sulphuric 
acid  liberate  carbon  monoxid;  with  dilute  sulphuric  acid 
hydrocyanic  acid,  HCN,  is  set  free.     (Danger) 

Hydroferricyanic  Acid,  H3Fe(CN)6,  and  Fenicyanids 
(the  Ferricyanid  Ion) 

Use  a  i-per-cent.  solution  of  KsFe(CN)6  for  the  reactions. 

212.  The  ferric  ion  with  soluble  ferricyanids  gives  no  pre- 
cipitate but  turns  the  liquid  brown  (56). 

213.  The  ferrous  ion  produces  a  deep  blue  precipitate  with 
the  f erricyanid  ion  (49) . 

Sulphocyanic  Acid,  HSCN,  and  Sulphocyanates 
(the  Sulphocyanate  Ion) 

A  i-per-cent.  solution  of  KSCN  may  be  used  for  testing. 

214.  The  ferric  ion  forms  a  deep  red  liquid,  but  no  precipi- 
tate, with  sulphocyanates  (57) .     It  disappears  upon  the  addi- 
tion of  mercuric  chlorid. 


ACIDS  (ANIONS)  107 

Hypochlorous  Acid,  HC1O,  and  Hypochlorites 
(the  Hypochlorite  Ion) 

A  solution  of  NaCIO  may  be  used  for  the  reactions. 

215.  Acidifying  hypochlorites  with  hydrochloric  acid  sets 
hypochlorous  acid  or  chlorin  free.     This  can  be  recognized, 
if  sufficient  is  present,  by  its  color,  odor,  and  bleaching  moist 
litmus-paper. 

216.  Soluble  hypochlorites  bleach  indigo  solution  —  more 
rapidly  after  acidifying. 

217.  Hypochlorites  do  not  bleach  acid  solutions  of  potas- 
sium permanganate. 

218.  The  lead  ion  forms,  first,  a  white  precipitate  which 
gradually  becomes  reddish,  then  brown  from  the  formation 
of  the  dioxid,  Pb02,  for  example, 

9NaC10+Pb(C2H302)2  =  Pb02+9NaCl+4C02+3H20. 
Nitrous  Acid,  HNO2,  and  Nitrites  (the  Nitrite  Ion) 

NaN02  can  be  used  for  the  reactions. 

219.  By  acidifying  solutions  of  nitrites,  nitrous  acid  is  set 
free.     It  is  extremely  unstable  and  is  not  used  in  the  free  state. 
From  the  solid  salts  or  concentrated  solutions  a  brown  mix- 
ture of  the  oxids  of  nitrogen  appears  when  acid  is  added. 


220.  The  free  acid  or  acidified  solutions  of  its  salts  decolor- 
ize indigo  solution  upon  heating. 

221.  Nitrites  when    acidified  will   decolorize  potassium 
permanganate. 

222.  From  a  few  drops  of  an  iodid  solution  acidified  with 
hydrochloric  acid,  nitrites  set  iodin  free.     Starch  solution 
then  gives  a  blue  color  with  the  iodin. 

2NaN02+2KI+4HCl=2KCl+2NaCl+N202+2H20+I2, 
or 


Ferric  salts  and  a  few  other  compounds,  notably  iodic  acid, 
also  decompose  an  iodid  in  a  similar  manner. 


108  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

• 

223.  The  silver  ion  forms  silver  nitrite,  AgN02,  which  is 
somewhat  soluble  in  water,  hence  does  not  appear  in  dilute 
solutions. 

Other  reactions  for  minute  amounts  of  nitrites  are  given 
under  Water  Analysis. 

Hydrochloric  Acid,  HC1,  and  Chlorids  (the  Chlorid  Ion) 

Hydrochloric  acid  is  a  colorless  gas  with  a  suffocating  odor, 
and  very  soluble  in  water,  in  which  form  it  is  commonly  used. 
Most  of  its  salts  are  also  soluble  in  water. 

224.  Addition  of  concentrated  sulphuric  acid  to  the  dry 
salts  produces  an  evolution  of  hydrochloric  acid  gas.     This 
has  a  strong  odor  and  acid  reaction.     If  manganese  dioxid, 
potassium  chlorate,  or  other  oxidizing  agent  is  also  present 
chlorin  is  set  free,  recognizable  by  its  color  and  odor. 

04  =  Na2S04+2HCl. 
=  MnCl2+2H20+Cl2. 

A  i-per-cent.  solution  of  NaCl  will  give  the  other  reactions. 

225.  The  silver  ion  precipitates  white,  curdy  silver  chlorid, 
AgCl,  which,  in  very  dilute  solutions,  remains   for  a  long 
time  in  suspension.     It  is  insoluble  in  nitric  acid,  but  easily 
dissolves  in  ammonium  hydroxid  and  can  be  reprecipitated 
from  this  solution  by  nitric  acid.     By  exposure  to  sunlight 
the  white  compound  turns  violet,  then  nearly  black. 

226.  The  lead  ion  forms  a  white  precipitate  of  lead  chlorid, 
PbCl2,  which  dissolves  in  boiling  water  and,  when  this  cools, 
separates  again  as  white  prismatic  crystals. 

Hydrobromic  Acid,  HBr,  and  Bromids  (the  Bromid  Ion) 

The  free  acid  is  similar  to  hydrochloric  in  physical  as  well 
as  chemical  properties.  For  the  reactions  use  potassium 
bromid. 

227.  From  a  crystal  of  a  dry  bromid  concentrated Jsul- 
phuric  acid  liberates  bromin  upon  warming.     This  is  seen 


ACIDS  (ANIONS)  109 

as  a  reddish-yellow  gas,  most  plainly  by  looking  down  into 
the  test-tube.  The  addition  of  a  little  manganese  dioxid 
increases  the  evolution  of  the  gas.  Nitric  acid  also  sets  free 
bromin  from  most  bromids  when  it  is  heated;  if  it  is  in  a  solu- 
tion it  colors  the  liquid  yellow.  From  a  solution  of  a  bromid, 
chlorin  water  liberates  free  bromin  and  this,  when  shaken 
with  a  drop  of  chloroform,  dissolves  giving  the  chloroform  a 
brown  color. 

228.  The  silver  ion  precipitates  from  solutions  of  bromids, 
even  when  dilute,  yellowish-white,  curdy  silver  bromid,  AgBr. 
It  does  not  dissolve  in  nitric  acid  and  is  soluble  with  some 
difficulty  in  ammonium  hydroxid. 

Hydriodic  Acid,  HI,  and  lodids  (the  lodid  Ion) 

The  acid  resembles,   physically  and  chemically,   hydro- 
chloric and  hydrobromic  acids. 

229.  A  small  crystal  of  the  dry  salt  is  decomposed  by  heat- 
ing with  concentrated  sulphuric  acid,  iodin  being  evolved  and 
volatilized  as  a  deep  purple  vapor.     Yellow  nitric  acid  pro- 
duces the  same  effect  when  boiled  with  the  salt  even  in  com- 
paratively dilute  solutions.     A  very  slight  amount  of  this 
vapor  turns  to  a  deep  bluish-black  a  paper  moistened  or  sized 
with  starch  solution. 

A  i-per-cent.  solution  of  potassium  iodid  may  be  used  for 
the  wet  reactions. 

230.  From  a  solution  of  an  iodid,  bromin  water  or  chlorin 
water  sets  free  iodin  which  gives  a  dark  blue  color  to  a  solu- 
tion of  starch.     An  excess  of  the  chlorin  water  destroys  the 
color.     lodids  alone  do  not  color  the  starch. 

231.  The  silver  ion  precipitates  the  iodid  ion  as  yellowish 
silver  iodid,  Agl,  insoluble  in  nitric  acid  and  nearly  insoluble 
in  ammonium  hydroxid. 

Hydrocyanic  Acid,  HCN,  and  Cyanids  (the  Cyanid  Ion) 

The  free  acid  is  a  colorless  liquid,  very  volatile,  with  the 
odor  of  bitter  almonds,  and  extremely  poisonous.     The  gas 


110  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

t 

burns  with  a  violet  flame.  The  cyanids  of  the  alkalies  and 
alkaline  earths  are  soluble  in  water.  Weak  acids,  including 
carbonic,  liberate  the  acid  from  these,  but  not  from  all 
cyanids.  The  insoluble  cyanids  of  the  heavy  metals  form 
soluble  cyanids  with  those  of  the,  alkalies,  consequently 
potassium  cyanid  dissolves  most  of  the  former  class. 
Use  for  the  reactions  a  i-per-cent.  solution  of  KCN. 

232.  The  silver  ion  forms  a  white  precipitate  of   silver 
cyanid,  AgCN,  which  is  insoluble  in  nitric  acid  and   but 
slightly  so  in  ammonium  hydroxid,  but  is  soluble  in  KCN, 
forming  KAg(CN)2,  the  potassium  salt  of  the  complex  ion 
Ag(CN),'. 

233.  A  few  drops  of  ferrous  sulphate  solution  with  a  solu- 
tion of  a  cyanid  is  converted  by  warming,  after  making 
alkaline  with  sodium  hydroxid,  into  the  ferrocyanid. 


That  is,  the  ferrous  ion  unites  with  the  cyanid  ion  to  form 
the  complex  ferrocyanid  ion. 

6CN'+Fe=Fe(CN)fl"". 

Upon  acidifying  with  hydrochloric  acid  and  adding  a  few 
drops  of  ferric  chlorid  (the  latter  is  unnecessary  if  the  ferrous 
sulphate  contained  any  ferric  salt),  a  deep  blue  color  ap- 
pears (55). 

234.  Acidification  of  solutions  of  cyanids  evolves  the  free 
hydrocyanic   acid,   recognizable   from  its   odor.     From  its 
poisonous  nature  it  is  dangerous  to  inhale  more  than  a  very 
little.     Concentrated  sulphuric  acid  decomposes  the  insol- 
uble cyanids  in  the  same  manner. 

Nitric  Acid,  HNO3,  and  Nitrates  (the  Nitrate  Ion) 

235.  The  free  acid  is  a  colorless  liquid,  very  strongly  cor- 
rosive.    It  colors  protein  compounds,  like  horn,    feathers, 
quill  toothpick's,  etc.,  a  bright  yellow. 


ACIDS  (ANIONS)  in 

Almost  all  nitrates  are  soluble  in  water  and  all  are  decom- 
posed with  the  evolution  of  oxygen  when  heated  to  a  red  heat. 
KNOa  can  be  used  for  the  reactions. 

236.  At  high  temperatures  they  give  up  their  oxygen  to 
oxidizable  substances  like  charcoal,  causing  deflagration.     A 
few  crystals  dropped  on  a  piece  of  red-hot   charcoal   will 
illustrate  this. 

237.  In  a  narrow  test-tube  place  3-4  c.c.  of  a  solution  of  a 
nitrate,  then  holding  it  in  a  slanting  position  slowly  pour  in 
as  much  concentrated  sulphuric  acid.     The  two  liquids  do 
not  mix.     If  much  heat  is  generated  cool  it  by  holding  the 
tube  in  running  water;  then  add  a  few  drops  of  a  concentrated 
solution  of  ferrous  sulphate.     A  brownish-purple  ring  forms 
at  the  junction  of  the  liquids.     With  minute  amounts  of  the 
acid  this  appears  only  slowly. 

238.  Nitric  acid  heated  with  metallic  copper  (or  nitrates 
with  this  and  in  addition  a  few  drops  of  concentrated  sul- 
phuric acid)  liberates  nitric  oxid  which  produces  a  brownish 
gas  in  the  upper  part  of  the  test-tube. 

239.  Nitric  acid  or  a  solution  of  a  nitrate  acidified  with 
hydrochloric  and  will  decolorize  indigo   on  warming.     It 
should,  however,  be  remembered  that  chlorin  and  some  other 
oxidizing  agents  will  do  the  same. 

Other  more  sensitive  tests  for  minute  amounts  of  nitric 
acid  are  given  under  the  subject  of  Water  Analysis. 

Chloric  Acid,  HC1O3,  and  Chlorates  (the  Chlorate  Ion) 

The  free  acid  is  not  used.  Its  salts  are  active  oxidizing 
agents  yielding  oxygen  more  easily  than  the  nitrates.  They 
deflagrate  on  red-hot  charcoal. 

Use  KClOa  for  the  reactions. 

240.  Chlorates  give  no  precipitate  with  the  silver  ion,  but 
when  heated  on  platinum  foil  they  lose  all  their  oxygen,  be- 


„, 


INTRODUCTION  TO   CHEMICAL  ANALYSIS 


coming  converted  to  chlorids  which  then  react  with  the  above 
reagent  (225). 

241.  Concentrated  sulphuric  acid  warmed  with  a  fragment 
of  a  dry  chlorate  liberates  chlorin  tetroxid,  C1C>2,  a  greenish- 
yellow  explosive  gas.     It  is  dangerous  to  use  large  amounts  of 
substance  or  to  heat  rapidly. 

242.  Hydrochloric  acid  will  evolve  chlorin  when  warmed 
with  a  chlorate,  as  is  proved  by  the  characteristic  color  and 
odor. 

KC10,+6HC1  =  KC1+3H20+6C1. 

Some  oxids  of  chlorin  may  also  be  formed.  The  liquid 
then  destroys  the  blue  color  of  an  indigo  solution. 

Acetic  Acid,  CHsCO^,  and  Acetates  (the  Acetate  Ion) 

The  free  acid  is  a  colorless  liquid  with  a  sharp  odor  and 
sour  taste.  It  is  easily  volatile  and  its  vapor  is  combustible. 
Most  of  its  salts  are  decomposed  by  heating  to  a  red  heat,  the 
residue  being  black.  But  few  acetates  are  insoluble. 

A  2-per-cent.  solution  of  CH3CO2Na  may  be  used  for  the 
reactions. 

243.  A  dry  acetate  heated  with  dilute  sulphuric  acid,  or 
a   solution  of   an   acetate   heated   with   concentrated   acid 
evolves  acetic  acid  which  distils  with  the  characteristic  odor. 

244.  If  in  the  above  test  a  cubic  centimeter  of  alcohol  is 
added  before  heating,  acetic  ether  (ethyl  acetate)  is  produced, 
and  this  can  be  identified  by  its  agreeable,  fruity  odor. 

245.  The  ferric  ion  imparts  a  deep  red  color  to  solutions 
of  neutral  acetates  or  to  acetic  acid  when  the  mixture  is  nearly 
neutralized  by  ammonium  hydroxid.     Boiling  this  precipi- 
tates the  ion.     Acidifying  with  hydrochloric  acid  turns  it 
yellow  (57). 


ACIDS  (ANIONS)  113 

Tartaric   Acid,  C^OiKCOzHK  and  Tartrates  (the  Tartrate  Ion) 

Tartaric  acid  is  a  white  crystalline  solid,  with  an  acid  taste, 
easily  soluble  in  water. 

Tartrates  prevent  the  precipitation  by  the  alkaline  hydrox- 
ids  of  copper,  iron,  and  some  other  bases  ordinarily  precipi- 
tated thus. 

C2H402(CO2)2KNa  may  be  used  for  the  reactions. 

246.  Dry  tartaric  acid  or  the  tartrates  are  charred  when 
warmed  with  concentrated  sulphuric  acid.     Both  it  and  its 
salts  are  decomposed  when  highly  heated  dry,  the  residue 
being  black  and  the  decomposition  attended  by  an  odor 
similar  to  that  of  burnt  sugar. 

247.  The  barium  ion  precipitates  from  solutions  of  tar- 
trates white  barium  tartrate,  .BaC^^e,  soluble  in  ammo- 
nium salts  or  hydrochloric  acid. 

248.  From  solutions  of  tartrates  the  silver  ion  precipitates 
white  silver  tartrate,  Ag2C4H406.     If  this  is  boiled  black 
metallic  silver  separates.     By  dropping  into  the  white  precipi- 
tate before  boiling  enough  ammonium  hydroxid  to  dissolve 
it,  then  slowly  warming  it  in  a  test-tube  that  has  been  thor- 
oughly cleaned  with  sodium  hydroxid  and  water,  the  silver 
deposits  as  a  mirror  on  the  tube.     The  heating  may  be 
accomplished  by  setting  the  tube  in  a  beaker  of  hot  water. 


Citric  Acid,  CgHgCKCO^H)^  and  Citrates  (the  Citrate  Ion) 

The  free  acid  is  a  crystalline,  colorless  solid  containing  one 
molecule  of  water  of  crystallization.  It  is  efflorescent, 
readily  soluble  in  water  and  has  an  acid  taste.  When  heated 
it  melts,  and  afterward  decomposes  with  blackening  and  the 
production  of  an  odor,  which,  however,  is  different  from  that 
from  tartaric  acid.  In  the  presence  of  citrates  the  alkaline 
hydroxids  fail  to  precipitate  iron  and  many  of  the  other  bases. 

KsCeHsOy  may  be  used  for  the  reactions. 


1 14  INTRODUCTION    TO   CHEMICAL   ANALYSIS 

249.  Citric  acid  and  its  salts  are  decomposed  by  hot  con- 
centrated sulphuric  acid  with  carbonization. 

250.  The  barium  ion  forms  a  white  precipitate  of  barium 
citrate,  Ba3(C6H5O7)2,  in  concentrated  solutions.     In  dilute 
solutions  it  may  not  appear  until  after  heating. 

251.  The  silver  ion  precipitates  white  silver  citrate,  Ag3C6- 
HsOy,  but  this  does  not  reduce  to  metallic  silver  when  it  is 
warmed  in  the  liquid. 

The  Identification  of  the  Acids  when  Only  One  is  Present 

As  a  preliminary  test  warm  2  c.c.  of  concentrated  sulphuric 
acid  nearly  to  boiling  in  a  test-tube,  then  drop  in  a  piece  of 
the  powdered  dry  substance  as  large  as  a  pea.  Avoid  heating 
to  the  boiling-point  of  the  sulphuric  acid.  Notice  whether 
there  is  an  evolution  of  gas  as  shown  by  the  formation  of  bub- 
bles in  the  hot  liquid. 

A.  A  Colorless  Gas  is  Evolved 

Chlorids  give  HC1  (odor). 

Cyanids  give  HCN  (odor).     (Danger!) 

Carbonates  give  CO2. 

Nitrates  give  HNOs  (odor). 

Sulphites  give  S02  (odor). 

Hypophosphites  give  S02  (odor). 

Sulphids  give  H2S  (odor),  although  some  sulphids  of  the 
heavy  metals  evolve  SO2  from  hot  sulphuric  acid. 

Ferrocyanids,  ferricyanids,  and  sulphocyanates  give  a 
mixture  of  gases,  sometimes  HCN.  (Danger!) 

Thiosulphates  give  SO2  (odor). 

Acetates  give  acetic  acid  (odor). 

Oxalates  give  CO  and  CO2. 

Some  salts  of  organic  acids  liberate  colorless  gases. 

B.  A  Colored  Gas  is  Evolved 

lodids  give  HI  and  iodin  (violet). 

Bromids  give  HBr  and  bromin  (reddish-brown). 


ACIDS  (ANIONS)  115 

Nitrates  give  HNO3  with  oxids  of  nitrogen  (brown  or  yel- 
lowish-brown) . 

Nitrites  give  oxids  of  nitrogen  (yellowish-brown). 

Chlorates  give  greenish-yellow  oxids  of  chlorin  (explosive). 

Hypochlorites  give  hypochlorous  acid  or  chlorin  (suffocat- 
ing odor). 

C.  No  Gas  is  Evolved  and  No  Blackening  is  Seen 

Sulphates. 

Phosphates. 

Borates. 

Chroma  tes. 

Arsenates. 

Arsenites. 

D.  The  Substance  is  Blackened 

Many  salts  of  organic  acids  as  well  as  some  organic  com- 
pounds which  are  not  salts  are  blackened  by  the  strong  sul- 
phuric acid  and  in  some  cases  colorless  gases  are  set  free. 

If  the  preliminary  test  with  sulphuric  acid  indicates  the 
acid  present,  confirm  it  by  the  identification  of  its  group,  then 
by  the  reactions  given  under  the  acid. 

In  case  the  preliminary  test  does  not  give  definite  results, 
dissolve  the  substance  in  water  if  possible,  if  not  in  a  very 
small  amount  of  nitric  acid.  Determine  the  class  to  which 
the  acid  belongs,  using  separate  portions  of  the  solution  for 
each  test.  To  identify  the  acid  after  this  has  been  done  com- 
pare its  reactions  with  those  of  all  the  acids  of  the  class. 

Although  it  is  possible  to  group  the  acids  in  accordance 
with  their  similarities,  the  groups  cannot  be  separated  from 
each  other  and  be  afterward  resolved  into  their  components 
as  easily  as  can  the  metals.  Such  an  analysis  is  often  accom- 
panied with  many  difficulties.  The  following  classification 
will  aid  in  the  identification  of  the  acid  radical  of  salts  when 


Il6  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

only  one  is  present  or,  if  there  are  several,  provided  they  are 
not  of  the  same  group. 

GROUP  I 

Precipitated  by  Ba  in  the  presence  of  HNO3. 

The  sulphate  ion,  SO/7  (sulphuric  acid  or  sulphates). 

GROUP  II 

Carefully  neutralize  the  solution  with  NH4OH  if  it  is  acid. 
If  the  cation  is  precipitated  filter  it  out1  and  test  the  nitrate 
for  the  anion  with  BaCl2.2  The  anions  of  this  group  give  a 
white  or  colored  precipitate,  soluble  in  HC1  or  HNO3.  They 
are  of  two  classes: 

Class  i. — The  substance  evolves  a  gas  with  hot  concen- 
trated H2S04. 

Carbonic  acid,  H2CO3  (carbonates).3 

Sulphurous  acid,  H2SO3  (sulphites).3 

Oxalic  acid,  H2C2O4  (oxalates). 

Citric  acid,  CsHsOy  (citrates). 

Class  2. — The  compound  does  not  evolve  a  gas  with  sul- 
phuric acid. 

Chromic  acid,  H2CrO4  (chroma tes). 

Phosphoric  acid,  H3PO4  (orthophosphates) . 

Arsenic  acid,  H3AsO4  (arsenates). 

Boric  acid,  H2BO3  (borates). 

GROUP  HI 

The  anions  of  this  group  give  precipitates  with  the  silver 
ion  in  the  presence  of  dilute  HNO3.  It  contains  four  classes. 

1  Salts  of  acids  in  group  II,  except  where  the  positive  constituent  is  an  alkali 
metal,  are  mostly  insoluble  in  neutral  or  alkaline  liquids.  Both  cation  and  anion 
will  therefore  in  such  cases  be  precipitated  when  the  solution  is  neutralized. 

2  If  the  cation  is  one  which  will  be  precipitated  by  the  chlorid  ion  (lead, 
silver,  or  mercurous),  Ba(No3)2  must  be  used  instead  of  BaCb- 

3  Carbonates  and  sulphites  which  are  insoluble  in  water  cannot  be  dissolved 
in  acid  and  the  solution  is  tested  with  a  barium  solution.    Why? 


ACIDS  (ANIONS)  117 

Class  i. — The  precipitate  with  AgNOs  turns  black  at  once 
or  on  warming.  Much  HNOs  prevents  this. 

The  thiosulphate  ion,  S2O3//  (thiosulphuric  acid,  or  thiosul- 
phates, sometimes  called  " hyposulphites").1 

The  sulphid  ion,  S"  (hydrosulphuric  acid,  or  sulphids, 
including  polysulphids). 

The  hypophosphite  ion,  PH2O2'  (hypophosphorous  acid, 
or  hypophosphites) . 

Class  2. — These  anions  give  a  red  or  blue  color  with  ferrous 
or  ferric  ions. 

The  sulphocyanate  ion,  SCN'  (sulphocyanic  acid  or  sul- 
phocyanates  or  "sulphocyanids"). 

The  hydroferrocyanid  ion,  Fe(CN)6r///  (hydroferrocyanic 
acid,  or  f errocyanids) . 

The  hydroferricyanid  ion,  Fe(CN)6///  (hydroferricyanic 
acid,  or  f erricyanids) . 

Class  3. — These  anions  bleach  indigo  solution  at  once  or 
on  warming. 

The  hypochlorite  ion,  CIO'  (hypochlorous  acid,  or  hypo- 
chlo  rites). 

The  nitrite  ion,  NCV  (nitrous  acid,  or  nitrites). 

Class  4. — Belong  to  none  of  the  first  three  classes. 

The  chlorid  ion,  Cl'  (hydrochloric  acid,  or  chlorids). 

The  bromid  ion,  Br'  (hydrobromic  acid,  or  bromids). 

The  iodid  ion,  V  (hydriodic  acid,  or  iodids). 

The  cyanid  ion,  CN'  (hydrocyanic  acid,  or  cyanids). 

GROUP  IV 

Give  no  precipitates  with  barium  or  silver  ions  in  the 
presence  of  nitric  acid. 

The  nitrate  ion,  N03'  (nitric  acid,  or  nitrates). 
The  chlorate  ion,  CKV  (chloric  acid,  or  chlorates). 

1  Since  silver  thiosulphate  is  soluble  in  an  excess  of  thiosulphates  it  may 
not  appear  if  too  concentrated  a  solution  of  the  thiosulphates  is  used. 


Il8  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

The  acetate  ion,  C2H3(V  (acetic  acid,  or  acetates). 
The  tartrate  ion,  C4H406//  (tartaric  acid,  or  tartrates). 

Practical  Exercises  on  the  Acids  and  their  Salts 

Determine  the  composition  of  the  unknown  compounds 
furnished  by  the  instructors  looking  up  their  physical  and 
chemical  properties,  their  manufacture,  their  uses,  etc.  If  a 
negative  result  is  obtained  in  the  examination  for  a  metal  the 
substance  may  be  an  uncombined  acid.  If  the  examination 
fails  to  show  any  acid  it  may  be  an  oxid  or  a  hydroxid.  The 
hydroxids  of  the  alkalies  and  alkaline  earths  form  alkaline 
solutions  in  water;  the  others  are  insoluble.  All  but  the 
hydroxids  of  the  alkalies  are  decomposed  to  water  and  an 
oxid  when  heated  in  a  dry  tube.  The  oxids  do  not  yield  water 
on  heating.  All  are  insoluble  in  water  except  those  of 
the  alkalies  and  alkaline  earths  which  dissolve  as  hydroxids. 
The  peroxids  and  a  few  others,  like  those  of  mercury,  give 
off  oxygen  when  heated. 

Questions  for  Further  Study  on  the  Reactions  of  the 

Acids 

Where  is  any  practical  use  made  of  the  affinity  of  concen- 
trated sulphuric  acid  for  water  ?  In  diluting  the  acid  which  is 
the  safer,  to  add  the  acid  to  the  water  or  to  add  the  water  to  the 
acid  ?  Will  the  concentrated  acid  char  such  an  organic  com- 
pound as  alcohol  by  removing  the  elements  of  water?  What 
gas  is  evolved  when  a  metal  dissolves  in  sulphuric  acid? 
Does  it  make  any  difference  whether  the  acid  is  concentrated 
or  dilute?  Is  it  a  compound  of  sulphuric  acid  which  is  pre- 
cipitated from  the  aromatic  acid  when  it  is  diluted  with  water  ? 
What  cations  are  precipitated  by  the  sulphate  ion?  Why  is 
sulphuric  acid  selected  for  making  the  preliminary  test? 


ACIDS  (ANIONS)  119 

Why  is  it  not  in  .all  cases  the  weaker  acid  which  is  thus 
liberated? 

Why  do  the  sulphites  so  often  contain  sulphates?  What 
effect  do  sulphites  have  on  solutions  of  salts  of  which  there 
are  two  classes,  the  ous  and  ic,  like  the  mercurous  and  mer- 
curic, ferrous  and  ferric,  arsenous  and  arsenic?  Will  other 
acids  than  sulphuric  set  free  the  sulphurous  acid  from  its 
salts?  What  is  the  chemical  action  of  sulphurous  acid  upon 
potassium  permanganate  or  the  chromates?  What  is  the 
chemical  action  by  which  hydrogen  sulphid  can  be  produced 
from  sulphurous  acid  and  its  salts,  and  what  is  the  active 
agent? 

With  solutions  of  what  metals  are  soluble  carbonates  in- 
compatible because  of  the  precipitation  of  the  former?  Are 
carbonates  ever  decomposed  by  salts  which- have  an  acid  reac- 
tion or  only  by  free  acids?  What  compound  is  formed  when 
carbon  dioxid  acts  upon  lime-water?  Wrhat  would  be  the 
effect  of  a  large  excess  of  the  gas?  How  are  the  bicarbonates 
produced  and  why  are  they  so  called?  What  application  is 
made  of  the  easy  decomposition  of  the  bicarbonates? 

What  is  the  physiological  action  of  oxalic  acid  or  the  soluble 
oxalates  ?  Are  many  of  the  other  organic  acids  volatilized  by 
heat  without  discoloration?  What  is  meant  by  " reducing 
power"?  After  barium  oxalate  has  been  dissolved  in  hydro- 
chloric acid,  why  is  it  precipitated  unchanged  by  adding  am- 
monium hydroxid  ? 

What  would  be  the  objection  to  triturating  a  dichromate 
with  tannic  acid  or  sugar?  Why  does  a  solution  of  a  dichro- 
mate turn  from  red  to  yellow  when  it  is  made  alkaline?  For 
what  is  chromic  anhydrid  used  ?  What  is  the  difference  be- 
tween a  salt  of  chromic  acid  and  a  chromium  salt? 

How  are  the  two  phosphoric  acids  made?  Will  soluble 
phosphates  be  incompatible  with  other  cations  than  iron, 
magnesium,  and  calcium?  The  phosphates  of  which  metals 


I2O  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

are  soluble  in  water?  Will  acids  and  alkalies  increase  or  de- 
crease the  solubility  of  phosphates  ?  What  is  the  commercial 
name  of  metaphosphoric  acid?  Which  phosphoric  acid  is 
official  in  the  U.S.P.  ?  How  much  of  the  hydrogen  of  ortho- 
phosphoric  acid  can  be  replaced  by  metals  to  form  salts? 
How  much  in  hypophosphorous  acid?  How  can  a  mineral 
acid  be  removed  from  a  solution  by  adding  sodium  acetate 
(190)  ?  How  do  salts  of  phosphoric  acid  formed  by  replacing 
varying  amounts  of  hydrogen  differ  from  each  other?  What 
use  is  made  of  the  zinc  compounds  of  metaphosphoric  acid? 
What  property  of  the  hypophosphites  accounts  for  their  de- 
colorizing a  permanganate  solution?  What  is  the  chemical 
change  which  occurs  when  hypochlorites  decolorize  indigo  so- 
lution? What  is  the  change  in  nitrites  when  they  decolorize 
a  permanganate  solution?  Is  their  chemical  action  always  of 
this  kind  when  mixed  with  other  compounds?  Is  the  chem- 
ical nature  of  the  organic  nitrites  used  medicinally  the  same 
as  of  the  inorganic  salts?  Is  the  chemical  activity  of  nitrites 
different  in  acidified  solutions  from  that  of  neutral  or  alkaline 
solutions? 

What  metallic  salts  of  the  halogens  are  insoluble  in  cold 
water?  What  is  the  effect  of  strong  oxidizing  agents  on 
hydrochloric  acid?-  How  do  the  compounds  of  chlorin, 
bromin,  and  iodin  compare  with  one  another  in  stability  as 
shown  in  their  decomposition  by  the  free  halogens  and  by 
other  agents  ?  What  is  the  action  of  their  salts  on  the  alka- 
loids? What  substances  liberate  iodin  from  iodids  and  are 
therefore  incompatible  with  these? 

In  what  combination  does  hydrocyanic  acid  occur  in  plants 
and  is  it  always  poisonous  in  such  form?  To  what  extent 
does  it  ionize  ?  In  what  plants  is  it  found  ?  How  stable  are 
solutions  of  the  acid  or  its  salts  ?  Are  any  of  the  salts  poison- 
ous? What  are  the  most  common  compounds  which  may  be 
oxidized  by  nitric  acid  ?  Does  it  ever  cause  explosions  thus  or 


ACIDS  (ANIONS)  121 

form  explosive  compounds?  What  is  its  action  upon  alka- 
loids? What  is  the  origin  of  the  yellow  color  frequently  seen 
in  nitric  acid ?  Are  the  nitrates  chemically  similar  to  the  free 
acid?  How  do  the  properties  of  the  chlorates  compare  with 
those  of  the  nitrates?  With  what  class  of  substances  are  dry 
chlorates  incompatible  and  why? 

What  change  is  produced  upon  the  tartrates  by  oxidizing 
agents  and  what  is  the  effect  upon  the  latter?  Where  are  tar- 
trates and  citrates  used  to  prevent  the  precipitation  of  cations 
by  the  alkalies? 


122      -  INTRODUCTION   TO   CHEMICAL   ANALYSIS 


CHAPTER  III 

ORGANIC   COMPOUNDS 
Benzin  and  Petroleum  Ether 

THESE  are  mixtures  of  hydrocarbons  of  the  marsh-gas 
series,  boiling  beween  55°  and  75°.  They  do  not  crystallize 
by  cooling  to  o°,  and  are  insoluble  in  water. 

252.  If  dropped  into  a  cooled  mixture  of  one  part  concen- 
trated sulphuric  acid  and  three  parts  fuming  nitric  acid  there 
is  no  discoloration  or  odor  of  bitter  almonds. 

253.  lodin  dissolves  in  either  to  a  violet  solution. 

Benzene  (Benzol),  CeH6 

A  colorless  liquid  with  a  characteristic  odor,  boiling  at 
about  80°,  lighter  than  water  and  almost  insoluble  in  it.  At 
zero  it  solidifies  to  a  crystalline  mass. 

254.  Cool  a  mixture  of  concentrated  sulphuric  and  nitric 
acids,  5  c.c.  of  each,  as  with  benzin  (252)  and  slowly  drop  in 
15-20  drops  of  benzene,  keeping  the  liquid  cool.     Nitroben- 
zene C6H5NO2,  is  formed,  a  yellowish  oil  with  the  bitter 
almond  odor. 

255.  lodin  dissolves  in  benzene  to  a  blood-red  solution. 

Chloroform,  CHC13 

A  colorless,  neutral  liquid  with  characteristic  odor,  having 
a  specific  gravity  of  1.5  and  a  boiling-point  of  about  61°.  It 
is  almost  insoluble  in  water. 

2  56.  A  strip  of  filter-paper  moistened  with  chloroform  burns 
with  a  green-edged,  yellow  flame. 

257.  A  drop  of  chloroform  should  not  be  rendered  milky 
by  a  drop  of  silver  nitrate  (225). 

258.  With  several  times  its  volume  of  alcohol  and  a  few 


ORGANIC   COMPOUNDS  123 

drops  of  sodium  hydroxid  it  is  decomposed  by  warming  into 
salts  of  hydrochloric  and  formic  acid. 

ChCl3+4NaOH  =  3NaCl+ 2HO2+HCO2Na. 

This  solution  gives  a  white  precipitate  with  silver  nitrate 
after  acidifying  with  nitric  acid  (225). 

259.  To  a  solution  of  potassium  hydroxid  in  alcohol  add  a 
drop   of   anilin   and  one  of  chloroform  and  warm  gently. 
Phenyl  carbylamin,  isonitril,  orisocyanid,  CeHsNC,  is  formed, 
recognizable  by  its  offensive  odor. 

260.  With  a  f  ew  crystals  of  beta-naphthol  and  a  few  drops  of 
sodium  hydroxid  chloroform  gives  a  blue  color  when  warmed 
slightly.     When  only  a  minute  quantity  is  present  the  blue 
may  be  transient. 

261.  In  a  10  c.c.  dry  test-tube  place  5  c.c.  of  chloroform  and 
add  without  agitation  5  c.c.  of  perfectly  clear  barium  hydroxid 
solution.     After  the  corked  tube  has  stood  six  hours  in  a  dark 
place  there  should  be  no  turbidity  at  the  line  of  contact  of  the 
liquids   (absence  of  decomposition  products  in  chloroform 
which  is  otherwise  pure). 

262.  Sodium  hydroxid  with  a  few  drops  of  copper  sulphate 
(Trommer's  reagent)  gives  a  reddish-yellow  precipitate  when 
warmed  with  a  drop  of  chloroform. 

263.  When  shaken  with  one-tenth  its  volume  of  concen- 
trated sulphuric  acid  and  allowed  to  stand,  pure  chloroform  is 
not  colored  yellow  or  brown. 

264.  When  pure  chloroform  has  been  thoroughly  shaken 
with  an  equal  volume  of  water  the  latter,  after  pouring  off, 
does  not  affect  litmus-paper,  is  unchanged  with  silver  ni- 
trate, and  gives  no  color  with  a  few  drops  of  potassium  iodid 
and  starch  solution  (230). 

lodofonn,  CHI3 

Yellow  crystals  with  a  characteristic  odor.     It  is  almost 
insoluble  in  water,  but  dissolves  in  alcohol  or  ether. 


124  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

265.  A  drop  of  the  alcoholic  solution  allowed  to  evaporate 
on  a  microscope  slide  leaves  the  iodoform  in  flat,  hexagonal 
or  star-shaped  crystals,  as  is  shown  by  the  microscope. 

266.  The  dry  substance  when  heated  in  a  test-tube  melts 
at  about  115°  to  a  brown  liquid  and  afterward  gives  off  the 
violet  vapors  of  iodin,  leaving  a  charred  mass.     The  latter  is 
completely  combustible  on  a  platinum  foil. 

267.  Water  which  has  been  shaken  with  one-fifth  its  weight 
of  pure  iodoform  and  filtered  is  colorless,  neutral,  has  no  bitter 
taste  and  is  unaffected  by  silver  nitrate. 

Ethyl  Alcohol,  C^OH 

A  colorless,  neutral  liquid  with  a  specific  gravity  of  0.79 
and  a  boiling-point  of  78°. 

268.  i  c.c.  of  alcohol  with  an  equal  volume  of  concen- 
trated sulphuric  acid  and  a  few  crystals  of  sodium  acetate 
gives,  when  heated  to  the  boiling-point,  ethyl  acetate,  C2H5- 
C2H3O2  (acetic  ether),  which  distills  with  a  pleasant,  fruity 
odor. 

269.  With  enough  sodium  hydroxid  to  make  it  alkaline, 
then  a  solution  of  iodin  in  potassium  iodid  until  it  has  a 
yellow  color,  followed  by  gentle  warming,  there  forms,  im- 
mediately or  after  standing,  a  light  yellow  precipitate  of  iodo- 
form (page  123).     This  reaction  is  a  very  sensitive  one.     It  is 
also  given  by  aceton,  acetic  aldehyd,  and  some  other  organic 
compounds,  but  not  by  methyl  alcohol. 

2  70.  If  enough  potassium  dichromate  is  added  to  color  the 
alcohol  light  yellow,  and  if  it  is  then  acidified  quite  strongly 
with  sulphuric  acid  it  is  reduced  to  a  chromic  salt  by  boiling 
and  this  colors  the  liquid  green.  This  is  not  characteristic 
of  alcohol,  as  other  reducing  agents  may  produce  the  same 
effect. 

271.  10  or  12  drops  of  alcohol  with  2  or  3  of  carbon bisul- 


ORGANIC   COMPOUNDS  125 

phid  and  a  little  concentrated  sodium  hydroxid  produce  a 
yellow  sodium  xanthogenate. 


This  with  an  aqueous  solution  of  ammonium  molybdate(i  :  10)  , 
after  acidifying  with  sulphuric  acid,  gives  a  red  color.  To 
obtain  it  with  dilute  alcoholic  solutions  a  few  drops  should  be 
well  mixed  with  one  of  carbon  disulphid  and  a  very  small  piece 
of  sodium  hydroxid  or  potassium  hydroxid;  after  standing  in 
an  evaporating  dish  at  the  ordinary  temperature  until  the 
carbon  disulphid  has  disappeared  a  drop  of  the  ammonium 
molybdate  is  to  be  added  and  enough  sulphuric  acid  to  acidify. 

272.  Pure  alcohol  (U.S.  P.)  should  mix  with  water,  ether, 
and  chloroform  without  cloudiness.     It  does  not  affect  the 
color  of  litmus-paper  and  leaves  no  residue  upon  evaporation  ; 
if  it  evaporates  spontaneously  from  a  paper  no  foreign  odor 
should  be  perceptible.     With  half  its  volume  of  potassium 
hydroxid  solution  it  should  not  at  once  become  dark  colored 
nor  should  silver  nitrate  give  more  than  a  faint  opalescence 
nor  more  than  a  faint  brownish  tint  when  standing  six  hours 
in  diffused  daylight. 

273.  To  test  for  methyl  alcohol  in  ethyl  alcohol  dilute  the 
liquid  if  necessary  with  water  so  that  it  contains  no  more  than 
10  per  cent,  alcohol;  make  a  spiral  7  mm.  thick  and  30  mm. 
long  by  winding  clean  copper  wire  around  a  glass  rod;  place 
10  c.c.  of  the  alcohol  in  a  30-40  c.c.  test-  tube  and  plunge  into  it 
the  spiral,  heated  to  redness,  repeating  the  operation  five  or  six 
times  and  keeping  the  liquid  cool  by  placing  the  tube  in  water. 
Filter  and  boil  gently  until  no  odor  of  acetaldehyd  remains. 
Cool  and  add  one  drop  of  a  5-per-cent.  solution  of  resorcin, 
then  pour  it  carefully  on  to  concentrated  sulphuric  acid  in  a 
narrow  test-tube  so  that  the  two  do  not  mix.     After  two  or 
three  minutes  rotate  the  tube  slowly.     More  than  2  per  cent. 
of  methyl  alcohol  gives  a  rose-red  ring  at  the  line  of  contact. 


126  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Amyl  Alcohol,  CsHnOH 

A  colorless  liquid  with  a  specific  gravity  of  0.82,  and  a  boil- 
ing-point of  131°.  It  dissolves  in  alcohol  but  not  in  water. 
It  has  a  peculiar  odor,  its  vapor  causing  coughing. 

274.  In  alcoholic  fluids  any  considerable  quantity  of  amyl 
alcohol  can  be  detected  by  means  of  its  odor  by  pouring  into  a 
large  beaker,  spreading  it  over  the  sides,  and  causing  the 
ethyl  alcohol  to  evaporate  quickly  by  swinging  through  the  air. 
Evaporation  on  a  water-bath  also  removes  the  water  and 
ethyl  alcohol  and  leaves  the  amyl  alcohol. 

275.  If  hydrochloric  acid  is  dropped  into  a  i-per-cent.  solu- 
tion of  methyl  violet  until  it  becomes  green,  then  amyl  alcohol 
is  mixed  with  it  in  a  porcelain  dish,  there  appear  on  the  top  of 
the  liquid  violet  drops. 

276.  Amyl  alcohol  dropped  into  i  c.c.  of  concentrated  sul- 
phuric acid  gives  a  red  color. 

Glycerin,  or  Glycerol  (C3H5OH)3 

A  neutral,  colorless,  odorless,  syrupy  liquid  of  a  specific 
gravity  of  1.25  to  1.27,  with  a  sweet  taste.  It  is  hygroscopic 
and  dissolves  in  water  or  alcohol. 

277.  A  borax  bead  dipped  in  a  solution  which  contains 
glycerol  colors  the  blue  flame  of  the  Bunsen  burner  green  (184) . 

278.  Glycerol  prevents  the  precipitation  of  copper,  lead, 
or  ferric  solutions  by  the  alkaline  hydroxids. 

279.  Five  drops  each  of  glycerol  and  concentrated  sulphuric 
acid  gently  warmed  with  as  much  resorcin  as  can  be  taken  up 
on  the  point  of  a  knife-blade  gives  first  a  crimson,  then  a  blood- 
red  color.     Too  high  heat  causes  charring.     When  diluted 
with  5  cc.  of  water  and  made  alkaline  with  ammonium  hydroxid 
the  yellow  solution  shows  a  green  fluorescence. 

280.  Pure  glycerol  will  give  no  precipitate  when  tested  by 
Trommer's  test  (262). 

281.  Pure  glycerol  when  warmed  gently  with  an  equal 


ORGANIC   COMPOUNDS  127 

volume  of  concentrated  sulphuric  acid  or  sodium  hydroxid  is 
not  colored  dark. 

282.  An  ammoniacal  solution  of  silver  nitrate  is  not  changed 
when  warmed  with  pure  glycerol,  but  upon  the  addition  to  it  of 
sodium  hydroxid  black  metallic  silver  separates  immediately. 

Phenol  (Carbolic  Acid),  C6H5OH) 

Pure  phenol  is  a  colorless,  crystalline  solid  with  a  character- 
istic odor.  It  melts  at  about  40°  and  dissolves  in  water  to  a 
neutral  solution.  The  melted  phenol  solidifies  upon  cooling, 
but  the  addition  of  a  little  water  prevents  this.  An  aqueous 
solution  may  be  used  for  testing. 

283.  Neutral  ferric  chlorid  solution,  not  in  excess,  produces 
a  deep  violet  color  which  is  changed  to  a  yellow  with  hydro- 
chloric acid. 

284.  Bromin  water  precipitates  first  white,  then  yellow 
tribromphenol,   CeH^BrsOH,  soluble  in  alkaline  hydroxids 
or  hydrochloric  acid. 

285.  A  solution  of  phenol  in  water  made  alkaline  with  a 
drop  of  ammonium  hydroxid  becomes  blue  when  warmed 
with  a  few  drops  of  bromin  water. 

286.  Five  c.c.  of  an  aqueous  solution  boiled  with  5  or  6 
drops  of  concentrated  nitric  acid  forms  a  brownish  or  yellow 
compound,  picric  acid,  CeH^OHXNC^a,  which  colors  a  large 
amount  of  water.     The  color  is  intensified  by  making  it 
alkaline  with  ammonium  hydroxid. 

287.  A  phenol  solution  with  a  drop  of  anilin,  then  an  excess 
of  a  solution  of  calcium  hypochlorite  gives  a  deep  blue  color. 
(Anilin  and  calcium  hypochlorite  give  a  red.) 

Creosote 

This  is  a  colorless  or  yellowish  liquid  with  a  smoky  odor, 
not  solidifying  in  the  cold  and  but  slightly  soluble  in  water. 

288.  Ferric  chlorid  colors  the  aqueous  solution  a  transient 


128  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

violet-blue  passing  into  a  grayish-green  and  brown,  lastly  a 
brown  precipitate.) 

289.  Bromin  water  gives  a  brownish  precipitate. 

Formaldehyd  (Formalin),  HCHO 

A  colorless  gas  with  a  very  irritating  odor.  It  dissolves  in 
water  and  is  commonly  to  be  obtained  in  the  form  of  a  40-per- 
cent, solution.  By  concentrating  the  aqueous  solution  the 
aldehyd  becomes  converted  into  the  solid  paraformaldehyd 
(HCHO)  3. 

A  i-per-cent.  solution  can  be  used  for  the  tests. 

290.  The  silver  ion  to  which  enough  ammonium  hydroxid 
has  been  added  to  dissolve  the  precipitate  first  formed  is  re- 
duced by  formaldehyd  to  metallic  silver  when  allowed  to 
stand  several  hours  in  a  dark  place,  the  metal  being  deposited 
as  a  mirror.     If  warmed  this  occurs  immediately. 

291.  Formaldehyd  reduces  a  solution  of  cupric  sulphate 
made  alkaline  with  sodium  hydroxid  with  the  formation  of 
a  reddish-yellow  precipitate  of  cuprous  oxid   (Trommer's 
reaction). 

292.  To  a  dilute  fuchsin  solution  acidified  with  sulphuric 
acid  add  just  enough  sodium  sulphite  to  decolorize  it.     For- 
maldehyd produces  a  purplish- violet  color  when  warmed  with 
this. 

293.  Formaldehyd  gives  a  yellowish  precipitate  when  it  is 
warmed   with  phenyl-hydrochlorid ;  this    is    the    hydrazon 
C6H5NHNCH2. 

294.  A  few  cubic  centimeters  of  a  dilute  formaldehyd  solu- 
tion with  about  50  mg.  of  resorcin  and  half  its  volume  of  50- 
per-cent.  sodium  hydroxid,  when  warmed,  is  turned  first 
yellow,  then  red.     It  is  a  very  sensitive  test. 

295.  Ammonium  hydroxid  converts  a  solution  of  formal- 
dehyd in  solution  into  hexa-methylen-tetra-amin,  which  re- 
mains as  a  solid  after  evaporation. 


ORGANIC   COMPOUNDS  129 

296.  A  solution  of  40-50  per  cent,  of  sodium  hydroxid  and 
5  per  cent,  of  resorcin  in  water  when  boiled  half  a  minute 
with  an  equal  volume  of  a  dilute  formaldehyd  solution  gives 
a  red  color.     Albuminous  compounds  prevent  the  reaction. 

297.  Fifteen  c.c.  of  a  very  dilute  solution  of  formaldehyd 
with  i  c.c.  of  a  dilute  solution  of  phenyl-hydrazin  hydro- 
chlorid  and  a  few  drops  of  a  freshly  prepared  solution  of 
sodium  nitroprussid  when  made  alkaline  with  sodium  hydroxid 
becomes  blue;  in  milk  a  grayish-green  color  results.     Chloro- 
form gives  a  similar  result,  but  only  in  a  much  stronger  solu- 
tion, where  it  can  usually  be  identified  by  its  odor. 

If  to  the  mixture  of  formaldehyd  and  the  phenyl-hydrazin 
salt  ferric  chlorid  is  added,  then  the  liquid  strongly  acidified 
with  concentrated  hydrochloric  acid  a  red  color  forms,  very 
slowly  changing  to  orange-yellow. 

If  milk  containing  formaldehyd  is  diluted  with  an  equal 
volume  of  water,  then  poured  on  to  concentrated  sulphuric 
acid  in  a  test-tube  so  as  to  produce  two  layers,  a  violet  ring 
appears  between  the  liquids.  This  is  Hehner's  test. 

These  three  last  tests  can  be  made  on  milk  directly  with- 
out distillation  and  all  are  very  sensitive. 

Chloral  Hydrate,  Hydrated  Chloral,  CC13CHO,H2O 

Colorless  crystals  with  a  sharp  odor,  soluble  in  water  or 
alcohol  to  a  neutral  solution,  melting  at  57°  and  boiling  at 

98°. 

298.  When    warmed    with    alkaline    substances    chloral 
hydrate  is  decomposed  with  the  formation  of  chloroform. 
The  latter  may  be  detected  by  its  odor  or  by  the  reactions 
for  chloroform  in  the  presence  of  an  alkali  (258,  259). 

299.  Chloral   hydrate   gives   no   precipitate   with  silver 
nitrate,  but  if  it  is  allowed  to  stand  for  a  time  with  dilute  sul- 
phuric  acid   and  zinc,    silver  nitrate  being  subsequently 
added,  silver  chlorid  is  precipitated. 


130  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

300.  From  an  ammoniacal  solution  of  silver  nitrate  chloral 
hydrate  precipitates  dark-brown  metallic  silver,  slowly  at  the 
ordinary  temperature,  more  rapidly  on  warming. 

301.  A  chloral  hydrate  solution  with  as  much  resorcin  as 
can  be  held  on  the  point  of  a  penknife  and  five  drops  of 
sodium  hydroxid  gives  an  intense  red  color  which  is  de- 
stroyed by  hydrochloric  acid. 

302.  Chloral  hydrate  with  ammonium  sulphid  gives  an 
orange  color  then  a  brown,  more  quickly  on  warming. 

303.  Pure  chloral  hydrate  should  not  give  the  iodoform 
reaction  (269). 

Benzole  Acid,  C6H5CO2H 

Colorless  needle-shaped  or  tabular  crystals  with  an  aromatic 
odor,  melting  at  121°,  soluble  with  difficulty  in  cold  water  but 
easily  in  warm  water,  soluble  in  alcohol,  ether,  and  chloroform. 

304.  When  heated  in  a  dry  tube  the  crystals  melt  and 
volatilize  without  blackening,  becoming  condensed  in  the 
crystalline  form  on  the  cooler  part  of  the  tube. 

305.  Benzoic  acid  dissolves  in  sulphuric  acid  without  char- 
ring. 

306.  The  lead  ion  precipitates  white  lead  benzoate,  Pb 
(C6H5CO2)2,  soluble  in  much  boiling  water  and  reprecipi- 
tated  on  cooling. 

307.  The  ferric  ion  precipitates  flesh-colored  basic  ferric 
benzoate,  Fe2(C7H5O2)3(OH)3.     Tartaric  or  citric  acid  pre- 
vents this. 

Salicylic  Acid,  C6H4OHCO2H 

Colorless,  odorless,  needle-shaped  crystals,  only  slightly 
soluble  in  cold  water,  more  soluble  in  hot  water.  The  normal 
salts  are  soluble  in  water,  the  basic  ones  less  so.  From  these 
solutions  the  free  acid  is  precipitated  by  mineral  acids. 

308.  When  slowly  heated  salicylic  acid  melts  at  156°  and 


ORGANIC    COMPOUNDS  131 

sublimes  as  white  crystals;  by  rapid  and  high  heating  it  de- 
composes into  carbon  dioxid  and  phenol,  the  latter  recogniz- 
able by  its  odor. 

309.  The  ferric  ion  gives  a  deep  bluish-purple  color.     Min- 
eral acids,  soluble  hydroxids  and  some  salts  may  prevent  the 
formation. 

310.  Salicylic  acid,  or  the  salicylates,  dissolved  in  i  c.c.  of 
methyl  alcohol,  when  i  c.c.  of  concentrated  sulphuric  is  added 
and   the  mixture  warmed,  will  produce  methyl  salicylate, 
C6H4(OH)(CO2CH3),    oil    of    wintergreen.     Ethyl    alcohol 
gives  ethyl  salicylate  in  the  same  manner,  having  a  very 
similar  odor. 

311.  Salicylic  acid  gives  similar  reactions  to  phenol  with 
boiling  nitric  acid  (286),  and  bromin  water  (284). 

312.  Concentrated  sulphuric  acid  makes  a  colorless  solu- 
tion with  salicylic   acid  when  cold,   turning  brown  upon 
warming. 

Meconic  Acid,  C5H2O3(CO2H)2 

313.  Meconic  acid  is  a  white  crystalline  solid,  soluble  with 
dimculty  in  cold  water,  but  readily  so  with  warming,  also 
soluble  in  alcohol. 

314.  The  ferric  ion  imparts  a  blood-red  color  to  a  solution 
of  meconic  acid.     This  is  not  discharged  by  boiling  (245),  by 
acidifying  with  hydrochloric  acid   (245),  nor  by  mercuric 
chlorid  (214).     Stannous  chlorid  decolorizes  it. 

315.  The  silver  ion  produces  a  yellowish  precipitate,  be- 
coming a  brighter  yellow  when  it  is  warmed. 

Tannic  Acid  or  Tannins 

There  are  a  number  of  tannins  with  similar  properties. 
They  are  yellowish  powders,  soluble  in  an  equal  weight  of 
water,  also  in  alcohol  or  glycerol,  and  turning  litmus-paper  red. 

316.  Solutions  of  tannic  acid,  when  made  alkaline,  rapidly 
absorb  oxygen,  becoming  red,  brown,  or  black. 


132  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

317.  Ferric  salts  produce  a  blue-black  color  or  precipitate 
which  is  prevented  or  destroyed  by  acids. 

318.  Lime-water  produces  a  bluish- white  precipitate,  be- 
coming darker  blue  as  the  amount  of  the  reagent  is  increased, 
and  finally  changing  to  a  pinkish  color. 

319.  Tannic  acid  reduces  a  permanganate  solution. 

320.  Solutions   of  albumin,   gelatin,   or   starch  are  pre- 
cipitated by  tannic  acid. 

321.  A  solution  of  iodin  in  potassium  iodid  to  which  enough 
potassium  cyanid  has  been  added  to  produce  a  colorless  liquid 
gives  a  red  color  with  tannic  acid.     An  excess  of  the  cyanid 
deepens  the  color.     It  changes  in  one  or  two  minutes  to  brown. 

322.  A    solution    of    antimony   potassium    tartrate   pre- 
cipitates a  tannic  acid  solution. 

323.  Tannic  acid  precipitates  most  alkaloids. 

Gallic  Acid,  CeH^OH^COsH 

A  white  or  yellowish,  crystalline  powder,  soluble  in  100  parts 
of  water,  also  in  alcohol  or  glycerol.     It  has  an  acid  reaction. 

324.  When  made  strongly  alkaline  solutions  of  gallic  acid 
become  yellow,  red,  and  brown  through  the  absorption  of 
oxygen.     If  the  alkali  is  not  in  excess  the  solution  slowly 
turns  green. 

325.  Ferric  salts  give  a  bluish-black  color  or  precipitate 
with  gallic  acid  solutions. 

326.  Lime-water  gives  a  bluish  precipitate,  as  the  reagent 
is  increased  darker  blue  in  color  by  reflected  light  and  a 
greenish  by  transmitted  light,  with  a  large  excess  changing  to 
a  pink  color. 

327.  A  permanganate  solution  is  reduced  by  gallic  acid. 

328.  Solutions  of  albumin,  gelatin,  or  starch  are  not  pre- 
cipitated by  gallic  acid. 

329.  Gallic  acid  does  not  precipitate  the  alkaloids. 

330.  With  a  solution  of  iodin  in  potassium  iodid  which  has 


ORGANIC    COMPOUNDS  133 

been  decolorized  by  potassium  cyanid,  gallic  acid  gives  a 
red  color,  changing  to  yellowish-brown.  The  latter  change  is 
slower  than  with  tannic  acid  (321). 

Starch,  (C6Hi0O5)x 

It  is  found  in  vegetable  substances  in  the  form  of  granules, 
those  of  one  species  of  vegetable  often  differing  sufficiently 
from  those  of  another  to  render  possible  determination  of 
their  source  by  microscopic  examination.  It  is  insoluble  in 
cold  water,  neutral  and  tasteless.  The  granules  are  destroyed 
by  heating  with  water,  the  starch  forming  a  semi-soluble 
paste.  This  may  be  used  for  the  tests. 

33  1  .  With  a  solution  of  iodin  starch  gives  a  dark  blue  color. 
This  is  decolorized  by  alkalies  or  by  heating.  In  the  latter 
case  when  the  liquid  cools  it  becomes  colored  again.  The 
dry  starch  is  colored  by  the  same  reagent. 

332.  When  boiled  some  time  with  a  dilute  mineral  acid 
starch  is  converted  into  glucose,  C6Hi2O6.  The  liquid  then 
forms  no  blue  compound  with  iodin. 


Glucose  or  Dextrose  (Grape  Sugar), 

This  is  to  be  obtained  as  a  dry  solid  or  as  a  colorless  syrup. 
Its  taste  is  less  sweet  than  that  of  cane-sugar.  It  is  readily 
soluble  in  water. 

333.  A  solution  of  glucose  when  made  alkaline  with  sodium 
hydroxid  and  into  which  copper  sulphate  has  .  then  been 
dropped  gives  on  warming  a  reddish-yellow  precipitate  of 
cuprous  oxid,  Cu2O  (Trommer's  reaction). 

334.  A  glucose  solution  when  heated  15  minutes  with  as 
much  phenyl-hydrazin  hydrochlorid  as  can  be  taken  up  on  the 
point  of  a  knife-blade  and  twice  as  much  crystallized  sodium 
acetate  will,  on  cooling,  deposit  bright  yellow  crystals  of 
phenyl-glucosazone,  best  observed  with  the  microscope. 

335.  Mixed  with  a  small  piece  of  compressed  yeast  and 


134  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

placed  in  an  inverted  test-tube,  standing  at  a  temperature  of 
40°  with  the  mouth  under  some  of  the  same  solution,  the 
glucose  ferments  to  carbon  dioxid,  which  rises  in  the  tube 
and  can  be  tested  by  Ikne-water  (171).  Alcohol  (269)  is 
formed  at  the  same  time. 

336.  When  war/med  with  glucose  an  ammoniacal  solution 
of  silver  nitrate  is  reduced  with  the  deposition  of  metallic 
silver. 

337.  Concentrated  sulphuric  acid  does  not  blacken  dry 
glucose  in  the  cold. 

Saccharose  (Cane-Sugar),  Ci2H22On 

Colorless,  odorless,  sweet  crystals,  neutral  in  reaction. 

338.  With  Trommer's  reagent  (333)  cane-sugar  gives  no 
precipitate  by  heating  to  the  boiling-point,  although  it  may 
do  so  by  long-continued  heating. 

339.  Boiling  with  dilute  mineral  acids  changes  cane-sugar 
into  glucose  which  may  be  detected  in  the  above  manner  after 
neutralizing  (333). 

340.  With  phenyl-hydrazin  saccharose  reacts  like  glucose 

(334). 

341.  An  aqueous  cane-sugar  solution  (1:10)  when  quickly 
heated  to  boiling  with  an  ammoniacal  solution  of  silver  nitrate 
then  allowed  to  cool  should  give  no  more  than  a  faint  colora- 
tion but  no  dark  precipitate  in  five  minutes. 

342.  Concentrated  sulphuric  acid  turns  dry  cane-sugar 
yellow,  brown,  and  at  last  black  without  the  application  of  heat. 

Lactose  (Milk-sugar),  CnH^On 

White,  hard  crystals  with  a  somewhat  gritty,  sweetish  taste, 
less  soluble  in  water  than  glucose  or  saccharose. 

343.  Lactose  gives  the  same  results  as  glucose  with  Trom- 
mer's reagent  (333). 


ORGANIC    COMPOUNDS  135 

344.  Yeast  does  not  produce  the  fermentation  of  lactose 

(335)- 

345.  Concentrated  sulphuric  acid  in  the  cold  does  not 

blacken  dry  milk-sugar. 

346.  Lactose  with  phenyl-hydrazin  (334)  and  Trommer's 
reagent  (333)  gives  similar  results  to  those  obtained  from 
glucose. 

Antipyrin  (Phenyl-di-methyl-pyrazolon),  CeHoCCHs^CsNoHO 

Colorless,  odorless  crystals  with  a  bitter  taste.  It  unites 
with  acids,  as  do  the  alkaloids,  to  form  salts,  but  has  no 
alkaline  reaction  to  litmus.  It  melts  at  113°  and  blackens  at 
higher  temperatures.  It  is  precipitated  by  most  of  the 
general  alkaloidal  reagents  (page  139). 

347.  Concentrated  sulphuric  acid  dissolves  antipyrin  to  a 
colorless  solution,  not  darkened  by  heating. 

348.  A  solution  of  antipyrin  in  water  when  made  strongly 
acid  with  nitric  acid  and  warmed  turns  yellow,  then  red 
(distinction  from  acetanilid  and  acetphenetidin). 

349.  The  solution  in  water  with  a  few  drops  of  fuming 
nitric  acid  becomes  emerald-green;  with  a  larger  amount  of 
the  acid  it  turns  red.     A  nitrite  added  to  a  slightly  acidified 
solution  of  antipyrin  also  gives  the  same  green  color. 

350.  The  aqueous  solution  of  antipyrin  is  colored  red  by  a 
drop  of  ferric  chlorid  solution  becoming  yellow  with  sulphuric 
acid. 

351.  Mercuric  chlorid  gives  a  white  precipitate,  readily 
soluble  on  warming  and  reprecipitating  as  the  solution  cools. 

Acetanilid  (Antifebrin),  C^NHCHgCO 

A  colorless,  odorless,  nearly  tasteless,  crystalline  sub- 
stance, soluble  with  difficulty  in  cold  water,  but  easily  soluble 
in  hot  water  or  alcohol.  It  melts  at  113°. 


136  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

352.  Acetanilid  dissolves  in  either  concentrated  sulphuric 
or  nitric  acid  without  discoloration  (347). 

353.  When  heated  with  sodium  hydroxid  acetanilid  is  de- 
composed with  a  formation  of  anilin  which  gives  its  character- 
istic odor.     If  then  a  few  drops  of  chloroform  are  added  and 
the  mixture  heated  it  gives  the  disagreeable  odor  of  isonitril 

(259)- 

354.  One  c.c.  of  hydrochloric  acid  warmed  with  a  little 
acetanilid  decomposes  the  latter  into  anilin  and  acetic  acid. 
If  a  drop  of  phenol  is  added,  and  a  little  calcium  hypochlorite 
solution,  the  liquid  becomes  reddish.     When  made  alkaline 
with  ammonium  hydroxid  this  changes  to  a  blue  (indophenol 
reaction). 

355.  The  ferric  ion  produces  no  color  with  acetanilid  be- 
yond the  yellow  of  the  reagent.     When  warmed  this  becomes 
darker. 

|356.  Warmed  with  i  c.c.  each  of  concentrated  sulphuric 
acid  and  alcohol  the  odor  of  ethyl  acetate  (268)  appears. 

Phenacetin  (Para-acet-phenetidin),  CaHsOCel^NHCHsCO 

Colorless,  odorless,  almost  tasteless  crystals,  soluble  in  hot 
water  or  alcohol,  almost  insoluble  in  cold  water.  It  melts  at 
135°  and  sublimes  at  a  higher  temperature  with  white  fumes. 

357.  Concentrated  sulphuric  acid  does  not  discolor  phen- 
acetin. 

358.  Warmed  with  concentrated  nitric  acid  a  yellow  or 
orange  color  appears. 

359.  Phenacetin  does  not  give  the  isonitril  reaction  (259) 
except  after  long  heating,  as  it  decomposes  much  more  slowly 
than  acetanilid  (353). 

i    360.  Phenacetin  gives  the  indophenol  reaction  like  acetani- 

lid  (354). 

361.  The  ferric  ion  gives  a  yellow  color  on  heating. 


ORGANIC;  COMPOUNDS  137 

362.  When  heated  with  concentrated  sulphuric  acid  and 
alcohol,  phenacetin  gives  ethyl  acetate  (356). 

363.  o.i  grm.  of  phenacetin  warmed  one  minute  with  i  c.c. 
of  concentrated  hydrochloric  acid,  then  diluted  to  10  c.c.,  on 
the  addition  of  three  drops  of  a  solution  of  chromic  anhydrid 
in  water  gives  a  ruby-red  color. 

»  Salol  (Phenyl  Salicylate),  C6H4OHCO2C6H5 

A  colorless,  crystalline  powder  with  an  aromatic  odor,  insolu- 
ble in  water  but  soluble  in  alcohol.  It  melts  at  42°;  in  hot 
water  it  goes  to  the  bottom  in  oily  drops. 

364.  When  warmed  with  water  and  a  few  drops  of  sodium 
hydroxid  it  dissolves  with  decomposition.     From  this  solu- 
tion hydrochloric  acid  precipitates  the  salicylic  acid  as  a 
white  solid  and  the  phenol  can  be  recognized  by  its  odor. 

365.  Equal  parts  of  salol  and  potassium  nitrate  mixed  with 
i  c.c.  of  concentrated  sulphuric  acid  form  a  bright  greenish- 
blue  liquid.     On  dilution  with  water  or  through  standing 
in  the  air  it  passes  into  red  and  fades.     Sulphuric  alone  pro- 
duces no  change. 

366.  From  the  alcoholic  solution  bromin  water  precipitates 
white  monobromsalol. 

367.  Oh  adding  to  an  alcoholic  solution  of  salol  a  few  drops 
of  a  very  dilute  (straw-yellow)  solution  of  ferric  chlorid  a 
violet  or  purple  color  appears;  if  the  salol  solution  is  dropped 
into  the  iron  solution  there  is  a  white  precipitate  but  no  color. 

368.  After  shaking  with  water  and  filtering,  the  filtrate 
should  give  no  result  with  the  ferric  ion  (283,  309),  the  barium 
ion  (164,  187),  nor  the  silver  ion  (225). 

Sulphonal  (Di-ethyl-sulphon-di-methyl-methan) , 

(C2H5S02)2C(CH3)2 

Colorless,  odorless,  tasteless  crystals,  soluble  with  diffi- 
culty in  cold  water,  but  readily  by  the  aid  of  heat,  also  in 
alcohol.  It  melts  at  126°. 


138  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

369.  On  platinum  foil  sulphonal  burns  with  a  yellow  flame 
and  the  odor  of  sulphurous  oxid. 

370.  If  sulphonal  is  mixed  with  ten  times  its  bulk  of  pow- 
dered charcoal  and  the  mixture  heated  to  redness  in  a  dry  tube  it 
gives  off  the  garlicky  odor  of  mercaptol.     Acetic  acid  and  sul- 
phurous oxid  are  formed  also.     If  instead  of  charcoal  the  sul- 
phonal is  fused  with  potassium  cyanid  mercaptol  is  formed  and 
also  potassium  sulphocyanate,  the  latter  being  found  in  the 
aqueous  solution  by  the  use  of  the  ferric  ion  after  nltering(2 14) . 

371.  Sulphonal    is    not    decomposed    by    mineral   acids, 
alkalies,  nor  the  halogens. 

The  Alkaloids 

The  alkaloids  are  a  class  of  vegetable  nitrogenous  com- 
pounds which  are  chemically  similar  to  ammonia.  Like  this 
the  free  alkaloids  are  of  a  basic  nature  and  they  form  salts  by 
uniting  with  acids  without  the  setting  free  of  hydrogen.  Only 
a  few  of  the  alkaloids  are  liquid,  most  of  them,  as  well  as  their 
salts,  being  crystallizable  solids.  Their  salts  are,  as  a  rule, 
much  more  easily  soluble  in  water  than  the  alkaloids  them- 
selves. They  are  decomposed  by  alkalies,  the  alkaloid  being 
set  free,  just  as  ammonia  is  freed  from  its  salts  by  the  same 
reagents.  When  the  solution  is  concentrated  this  process  often 
results  in  their  precipitation,  as  the  alkaloid  is  less  soluble 
than  its  salts.  There  is  a  great  difference  in  the  comparative 
solubility  of  the  alkaloids  in  many  of  the  organic  liquids  like 
chloroform,  benzene,  ether,  petroleum  ether,  amyl  alcohol, 
etc.,  and  the  methods  devised  for  their  separation  are  based 
upon  this.  They  all  have  a  bitter  taste,  as  do  their  salts. 
The  separation  and  identification  of  the  alkaloids  is  not  as 
simple  an  operation  as  that  of  the  metallic  compounds.  It 
should  be  borne  in  mind  that  a  very  small  amount  of  impurity 
may  conceal  or  to  a  great  degree  modify  the  reactions. 

There  are  a  number  of  reagents  which  will  precipitate  most 


ORGANIC    COMPOUNDS  139 

or  all  of  the  alkaloids  and  these  are  of  value  in  proving  the 
presence  or  absence  of  the  class,  although  the  results  are  often 
so  similar  that  they  cannot  be  used  for  the  identification  of 
the  individual  members.  Among  the  most  important  of 
these  alkaloidal  group  reagents  are: 

Tannic  acid,  which  precipitates  most  alkaloids,  as  well  as 
some  other  similar  substances,  as  white  or  yellowish,  floccu- 
lent  compounds.  They  are  often  soluble  in  excess  of  the 
precipitant  or  in  other  acids. 

Picric  acid,  which  from  not  too  dilute  solutions  precipitates 
yellow  compounds,  often  crystalline  in  form. 

Phosphomolybdic  acid  precipitates  the  alkaloids  and  similar 
nitrogenous  compounds  in  the  form  of  yellowish  or  brown- 
ish-yellow solids.  These  can  be  filtered  from  the  solution  and 
the  alkaloid  set  free  from  them  by  the  alkalies  and  their 
carbonates. 

Phosphotungstic  acid  acts  like  the  phosphomolybdic  in  most 
cases. 

Mercuric  potassium  iodid  precipitates  most  alkaloids  from 
solutions  of  their  sulphuric  or  hydrochloric  acid  salts  as  white 
or  yellow  compounds. 

lodin  in  potassium  iodid  forms  brown  precipitates  with 
alkaloidal  solutions. 

\Mercuric  chlorid,  platinic  chlorid,  or  gold  chlorid,  from  not 
too  dilute  solutions,  throw  down  wrhite  or  yellowish  precipi- 
tates. With  dilute  solutions  they  may  only  form  a  turbidity. 

In  order  to  identify  a  compound  as  an  alkaloid  it  must  not 
only  give  the  general  tests  for  the  class  but  be  characterized 
by  some  special  reaction  or  reactions.  In  the  special  reac- 
tions that  follow  it  may  be  assumed,  unless  it  is  otherwise 
stated,  that  the  alkaloid  undej  discussion  will,  in  the  main, 
give  the  general  reactions. 

In  the  reactions  of  the  alkaloids  and  in  others  where  it  often 
becomes  necessary  to  identify  minute  quantities  of  material 


140  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

it  is  convenient,  instead  of  test-tubes,  to  make  use  of  small, 
rather  flat  watch-crystals.  The  results  of  the  treatment  of  a 
drop  of  solution  with  one  of  the  reagent  are  plainly  percepti- 
ble if  the  watch-glass  stands  upon  a  suitably  colored  paper, 
and  the  precipitate  can  be  examined  with  the  microscope, 
placing  the  watch-glass  upon  the  stage.  It  if  is  desired  to 
make  the  test  upon  a  solid  compound  the  drops  of  solution 
on  the  watch-glass  may  be  evaporated  to  dryness  on  the 
steam-bath.  At  first  rather  large  amounts  of  the  alkaloids 
may  be  used,  but  after  some  skill  has  been  gained  this 
should  be  reduced  to  very  minute  quantities  in  order  to 
show  the  sensitiveness  of  the  reagents. 

Morphin,  Ci7Hi9NO3 

Morphin  and  its  sulphuric  and  hydrochloric  acid  salts  are 
white  compounds  crystallizing  in  needles  or  prisms. 

372.  Concentrated  nitric  acid  dissolves  morphin  or  its  salts 
with  a  blood-red  color  which  gradually  passes  into  a  yellow. 
The  latter  is  not  changed  to  a  violet  by  addition  of  stannous 
chlorid,  but  is  colored  reddish-brown.     Only  a  small  amount 
of  the  acid  should  be  used  (difference  from  quinin) . 

373.  Concentrated  sulphuric  acid  dissolves  morphin  to  a 
colorless  liquid.     If  to  this  a  trace  of  nitric  acid  is  added  there 
is  no  change  in  color,  but  if  it  has  stood  24  hours,  or  is  heated  to 
100°  for  half  an  hour,  the  nitric  acid  produces  a  red.     A  crystal 
of  potassium  permanganate  added  to  the  concentrated  sul- 
phuric acid  solution  should  produce  no  violet  or  purple  color. 

374.  Frohde's  reagent  gives  a  violet  solution,   gradually 
changing  to  brown,  green,  and  yellow. 

375.  A  neutral  solution  of  ferric  chlorid  with  morphin  or 
solutions  of  its  neutral  salts  gives  a  blue  color.     Only  a  small 
quantity  of  the  ferric  compound  can  be  used,  as  an  excess 
prevents  the  reaction.     The  color  is  destroyed  by  free  acids 
and  impurities   may  prevent  t  its   appearance.     The   ferric 


ORGANIC   COMPOUNDS  141 

chlorid,  which  ordinarily  has  an  acid  reaction,  must  be  pre- 
pared by  subliming  the  crystals  in  a  hard-glass  tube  after 
driving  off  the  water  and  acid  by  gentle  heating;  then  dis- 
solve the  sublimate  in  distilled  water.  The  solution  should 
be  nearly  or  quite  neutral.  Impurities  may  prevent  the  for- 
mation of  the  color.  The  correctness  of  the  reagent  must  be 
proved  by  its  ability  to  produce  the  blue  with  a  specimen  of 
pure  morphin  sulphate.  Some  ptomains  give  a  greenish 
color,  but  not  a  blue.  None  of  the  other  vegetable  alkaloids 
give  this  result. 

376.  Morphin  with  half  its  weight  of  sugar  gives  a  red 
color  to  a  drop  of  concentrated  sulphuric  acid. 

377.  Vanadium  sulphate  produces  a  reddish,  then  a  violet 
color. 

Narcotin,  C22 


Concentrated  sulphuric  acid  dissolves  narcotin  to  a  colorless 
liquid.  If  the  acid  contains  a  trace  of  nitric  acid,  as  happens 
often  with  the  ordinary  chemically  pure  acid,  the  solution  turns 
yellow  on  standing  and  red  by  warming.  The  red  is  easily 
obtained  by  dissolving  in  a  little  dilute  acid,  then  evaporating 
from  a  test-tube  very  slowly  by  means  of  a  small  flame.  After 
the  red  liquid  cools  a  trace  of  nitric  acid  changes  it  to  violet. 

378.  Concentrated   nitric   acid    dissolves   narcotin    to    a 
yellow  solution. 

379.  Frohde's  reagent  of  the  ordinary  strength  dissolves 
narcotin,  forming  a  green  solution.     If  stronger  (o.oi  gm. 
molybdate  in  i  c.c.  of  acid)  the  green  solution  soon  changes 
to  a  cherry-red. 

380.  Erdmann's  reagent  gives  a  red  solution. 

381.  Vanadium  sulphate  gives'  a  red  color  likewise. 

Strychnin,  C2iH22N2O2 

Strychnin  is  soluble  with  difficulty  in  water,  its  salts  much 
more  easily  so.  It  is  one  of  the  most  intensely  bitter  of  the 


142  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

alkaloids.  It  gives  reactions  in  very  dilute  solutions  with 
most  of  the  general  alkaloidal  reagents.  The  chlorids  of  gold 
and  platinum  do  not  react  in  very  dilute  solutions,  perhaps 
below  o.i  per  cent,  to  o.oi  per  cent. 

382.  Strychnin  forms  a  colorless  solution  with  concen- 
trated sulphuric  acid. 

383.  To  the  dry  substance  add  a  drop  of  concentrated  acid, 
about  one  part  water  to  five  parts  of  acid,  then  by  means  of  a 
glass  rod,  draw  through  the  solution  a  minute  crystal  of  potas- 
sium dichromate.     A  series  of  colors  results,  always  in  the 
same  order;  first  blue  which  quickly  becomes  violet,  then 
more  slowly  red,  pink,  and  yellow.     The  colors  are  character- 
istic of  strychnin  and,  when  the  substance  is  pure  will  be 
produced  by  as  small  an  amount  as  the  fifty-thousandth  of  a 
grain.     The  reaction  is,  however,  interfered  with  by  the  pres- 
ence of  a  number  of  organic  compounds,  including  morphin. 

384.  Vanadium  sulphate  in  concentrated  sulphuric  acid 
dissolves  the  solid  compounds  with  the  production,  first  of  a 
blue,  followed  by  a  violet  and  red  color.     If  it  is  then  diluted 
with  water  the  pink  remains  for  a  long  time. 

385.  Cerium  oxid  with  strychnin  in  concentrated  sulphuric 
acid  gives  the  same  colors  as  the  last  reagent.     It  is  said  to 
react  with  the  millionth  of  a  gramme.     To  make  the  reagent 
heat  cerium  oxalate  to  redness  to  form  the  oxid  and  dissolve 
this  in  thirty  times  its  weight  of  concentrated  sulphuric  acid. 

386.  Strychnin  gives  no  color  when  treated  with  concen- 
trated sulphuric  acid  and  a  molybdate,  nitrate,  or  nitric  acid. 

387.  Concentrated  nitric  acid  colors  strychnin  or  its  salts 
only  faintly  yellow,  if  at  all,  when  it  dissolves  them,  but  on 
standing  the  solution  becomes  a  darker   yellow  (difference 

from  brucin). 

Brucin, 


The  alkaloid  and  most  of  its  salts  are  crystalline. 

388.  Concentrated  nitric  acid  dissolves  brucin  to  a  deep  red 


ORGANIC    COMPOUNDS  143 

liquid,  which  on  standing  or  heating  becomes  yellow.  After 
this  has  occurred  dilute  the  solution  and  with  this  a  reducing 
agent,  like  stannous  chlorid,  will  give  a  violet.  An  excess  of 
nitric  acid  must  be  avoided  or  the  reaction  loses  in  sensitive- 
ness. The  violet  solution  when  made  alkaline  with  sodium 
hydroxid  changes  to  blue  or  green.  These  changes  are  pecu- 
liar to  brucin. 

389.  Concentrated  sulphuric  acid  dissolves  brucin  without 
producing  a  color. 

390.  Chlorin  water  forms  a  bright  red  solution,  the  color 
being  destroyed  by  an  excess  of  the  .reagent  and  turned 
brownish-yellow  by  ammonium  hydroxid. 

Atropin  (Daturin),  C17H23NO3 

391.  Atropin  imparts  a  color  to  neither  concentrated  nitric 
nor  sulphuric  acid  in  the  cold,  but  dissolves  to  a  clear  solution. 

392.  A  crystal  of  atropin  when  moistened  with  three  or 
four  drops  of  fuming  nitric  acid  leaves  a  yellowish  residue 
when  evaporated  to  dryness  on  a  steam-bath.     Let  a  drop  of 
a  solution  of  potassium  hydroxid  in  90-per-cent.  alcohol  run 
over  the  residue.     Where  they  come  in  contact  a  reddish- 
violet  color  is  produced. 

393.  If  a  milligramme  of  atropin  is  heated  in  a  dry  test- 
tube  until  vapors  begin  to  appear  and  then  i  c.c.  of  concen- 
trated sulphuric  acid  is  added  and  the  heating  continued,  the 
odor  of  flowers  is  observed.     By  cautiously  diluting  with  2  c.c. 
of  water  and  warming  further  it  may  be  made  more  distinct. 
On  addition  of  small  crystals  of  potassium  dichromate  to  the 
concentrated  acid  solution  the  odor  changes  to  that  of  oil  of 
bitter  almonds. 

394.  Platinic  chlorid  does  not  precipitate  atropin  or  its 
salts  (difference  from  most  alkaloids). 

395.  A  very  dilute  aqueous  solution  of  atropin  (i  :  100,000) 


144  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

when  applied  to  the  eye  causes  a  dilation  of  the  pupil  which 
lasts  for  a  long  time. 

Veratrin 

396.  Veratrin   dissolves   in   concentrated   sulphuric   acid 
which  is  free  from  nitrogen  compounds  to  a  colorless  solution. 
On  warming  it  changes  to  a  crimson-red.     In  the  ordinary 
chemically  pure  acid  it  makes  a  yellow  solution,  changing  to 
orange,  then  blood-red  and  in  half  an  hour,  to  a  carmine-red. 
The  addition  of  a  minute  quantity  of  nitric  acid  hastens  these 
changes. 

397.  When  warmed  a  few  minutes  in  a  small  test-tube  with 
i  c.c.  of  concentrated  hydrochloric  acid  the  latter  is  colored  a 
reddish- violet.     No  other  alkaloid  gives  this  reaction. 

398.  If  mixed  with  6  times  the  amount  of  cane-sugar,  then 
moistened  with  concentrated  sulphuric  acid  (only  a  little 
being  added  from  a  glass  rod,  not  enough  to  make  a  solution) 
veratrin  gives  a  dark  green  color,  becoming  purple  and  then 
blue  in  a  few  times.     From  the  action  of  the  acid  on  the 
sugar  it  finally  turns  brown  and  black. 

Cocain,  Ci7H2iNO4 

399.  Solutions  of  the  alkaloid  or  its  salts  in  water  cause 
numbness  when  applied  to  the  tongue  and  dilation  of  the 
pupil  when  applied  to  the  eye. 

400.  Concentrated  sulphuric  acid  dissolves  the  cocain  or 
its  salts  without  discoloration.     When  warmed  it  turns  brown, 
the  hydrochlorid  giving  the  odor  of  hydrochloric  acid.     The 
vapors  produced  by  heating  with  the  concentrated  sulphuric 
acid  deposit  benzoic  acid  on  the  sides  of  the  test-tube  on  cooling. 

401.  Concentrated  nitric  acid  gives  no  color  with  cocain 
in  the  cold. 

402.  With  i-per-cent.  solution  of  potassium  permanganate 
cocain  yields  a  precipitate  of  cocain  permanganate.     This 


ORGANIC   COMPOUNDS  145 

upon  standing  collects  in  radiating  clusters  of  thin  tabular 
crystals,  a  bright  reddish-purple  in  color,  very  plainly  to  be 
seen  with  a  low  power  of  the  microscope. 

403.  An  aqueous  solution  of  cocain  acidulated  with  dilute 
sulphuric  acid,  to  which  enough  potassium   permanganate 
solution  has  been  added  to  give  a  violet  color  does  not  lose 
this  after  standing  half  an  hour  at  the  ordinary  temperature. 

404.  If  a  little  cocain  hydrochlorid  is  powdered  with  as 
much  mercurous  chlorid  in  a  dry  porcelain  dish  the  mixture  is 
white,  but  it  becomes  gray  in  presence  of  a  little  moisture, 
even  by  being  breathed  upon. 

405.  Cocain  or  its  hydrochlorid  when  evaporated  to  dry- 
ness  on  the  steam-bath,  after  the  addition  of  i  c.c.  of  concen- 
trated nitric  acid,  leaves  a  colorless  residue.     With  a  few 
drops  of  alcoholic  potassium  hydroxid  it  develops  the  pleasing 
and  permanent  odor  of  benzoic  ethyl  ester. 

406.  Cocain  hydrochlorid  solution  with  a  few  drops  of 
potassium    dichromate    gives    a   yellow    precipitate  which 
quickly  disappears.     On  acidifying  the  solution  then  with 
hydrochloric  acid  an  orange-yellow  crystalline  precipitate  of 
cocain  chromate  appears  which  is  soluble  in  excess  of  the  acid. 

Quinin,  C2oH24N2O2 

407.  Quinin  or  its  salts  give  no  color  with  Erdmann's  re- 
agent,  Frohde's  reagent  or  with  concentrated  nitric  acid 
(difference  from  morphin). 

408.  Solutions  of  quinin,  as  well  as  many  salts  of  quinin, 
after  acidifying  with  sulphuric  acid,  give  a  bluish  fluorescence, 
.perceptible   in   very   dilute   solutions   and   a  characteristic 
reaction.     Concentrated  acid  dissolves  the  dry  alkaloid  to  a 
similar  fluid,  with  no  brown  or  black  color. 

If  the  solution  in  sulphuric  acid  is  exactly  neutralized  with 

ammonia,  one  drop  of  hydrogen  peroxid  and  one  of  copper 
10 


146  INTRODUCTION  TO   CHEMICAL  ANALYSIS 

sulphate  solution  added,  then  boiled,  an  intensely  red  color 
appears,  changing  slowly  to  blue,  and  finally  green. 

409.  Ten  c.c.  of  a  solution  of  a  quinin  salt  in  water  with  two 
drops   of  bromin  water  or  chlorin  water  and  enough  am- 
monium hydroxid  to  render  it  alkaline  gives  an  emerald-green 
color — a  characteristic  reaction. 

410.  If  to  an  alcoholic  solution  of  quinin  sulphate  there  is 
added  tincture  of  iodin,  and  the  mixture  is  warmed,  then 
allowed  to  stand  and  cool,  there  separates  a  precipitate  con- 
sisting of  clusters  of  crystals,  dark  green  by  transmitted  light 
and  with  a  metallic  luster  by  reflected  light.     The  compound 
is  called  herapathite,  and  is  characteristic  of  quinin. 

Questions  for  Further  Study  on  Organic  Compounds 

What  are  organic  substances  and  how  distinct  is  the  line 
between  these  and  the  inorganic?  Are  benzene  and  benzin 
easily  combustible?  Why?  What  is  the  nature  of  sub- 
stances of  which  the  vapors  form  explosive  mixtures  with  air? 
What  is  the  difference  in  arrangement  of  atoms  in  the  mole- 
cules of  the  above  compounds?  Under  what  circumstances 
is  iodoform  decomposed  by  air  or  light?  What  oxidizing 
agents  oxidize  alcohol  or  its  solutions?  What  is  produced  by 
the  action  of  strong  mineral  acids  upon  alcohol?  Why  do 
many  tinctures  give  precipitates  when  diluted  with  water? 
Are  there  any  which  do  not?  Why?  What  compounds, 
organic  and  inorganic,  are  soluble  in  alcohol?  In  what  bev- 
erages is  amyl  alcohol  found  and  under  what  name  ?  What  is 
the  value  of  glycerin  as  a  solvent?  Will  carbolic  acid  neutral- 
ize solutions  of  the  alkalies  ?  With  what  solids  does  crystal- 
lized phenol  form  a  liquid  mass  on  trituration  ?  For  what  is 
formaldehyd  most  extensively  used?  How  can  it  be  most 
readily  generated?  Where  does  meconic  acid  occur  and 
what  is  its  importance?  What  drugs  contain  a  large  proper- 


ORGANIC   COMPOUNDS  147 

tion  of  tannic  acid?  Are  the  aqueous  solutions  of  gallic  and 
tannic  acid  permanent  ?  What  would  be  the  effect  of  triturat- 
ing them  dry  with  strong  oxidizing  agents?  Is  tannic  acid 
incompatible  with  solutions  of  any  other  metals  except  anti- 
mony? Can  starch  be  changed  into  glucose  by  other  means 
than  by  mineral  acids?  What  is  the  relation  of  this  to  " liver 
starch"?  Does  grape-sugar  occur  elsewhere  than  in  the 
grape?  What  are  the  best  means  of  distinguishing  the  three 
kinds  of  sugar?  What  is  the  relation  of  the  alkaloids  to  the 
ptomaines?  Is  there  any  similarity  in  their  reactions? 
Which  are  the  liquid  alkaloids?  With  what  solids  does  anti- 
pyrin  give  a  liquid  mass  on  trituration?  What  is  the  differ- 
ence between  an  isonitril  and  the  cyanid  of  an  organic 
radical?  What  are  the  mercaptols?  Trituration  with  what 
solids *and  salol  gives  a  liquid  mass? 


PART  II 

VOLUMETRIC  ANALYSIS 


CHAPTER  I 

GENERAL    PRINCIPLES 

In  quantitative  chemical  analysis,  that  is,  the  determina- 
tion of  the  amounts  of  the  constituents  of  a  compound  or  mix- 
ture, two  methods  may  be  employed,  the  gravimetric,  where 
the  constituent  or  some  derivative  of  it  is  isolated  and 
weighed,  and  the  volumetric,  where  to  a  definite  quantity  of 
the  substance  there  is  added  of  a  dissolved  reagent  just  suffi- 
cient to  complete  some  chemical  change,  such  as  precipita- 
tion, neutralization  or  oxidation.  If  the  concentration  of  this 
reagent  is  known  we  can  calculate  from  the  volume  necessary 
for  the  reaction  the  weight  of  the.  compounds  acted  upon. 

For  example,  the  concentration  of  a  solution  of  silver  which 
has  been  dissolved  in  nitric  acid  may  be  ascertained  in  two 
ways.  We  know  that  when  the  metal  dissolves  it  forms  sil- 
ver nitrate,  of  which  the  formula  is  AgNO3.  Of  this  the 
atom  of  silver  weighs  107.88,  of  nitrogen  14.01,  and  the  three 
atoms  of  oxygen  48.0,  or  a  weight  of  169.89  for  the  molecule. 
If  hydrochloric  acid  is  added  to  this  solution  silver  chlorid  is 
produced. 

AgN03  +  HC1  =  AgCl  +  HN03. 
169.89        36.468     144.34     63.018 

That  is,  one  molecule  of  hydrochloric  acid  having  a  weight 
of  36.468  will  convert  169.89  parts  of  silver  nitrate  into  144.34 
parts  of  silver  chlorid  and  63.018  parts  of  nitric  acid.  The 

148 


VOLUMETRIC   ANALYSIS  149 

silver  chlorid,  which  forms  a  precipitate,  can  be  filtered  from 
the  liquid,  washed,  dried,  and  weighed.  Since  of  every 
144.34  parts  of  the  chlorid  there  are  present  35.46  of  chlorin 
and  107.88  of  silver  a  simple  proportion  will  give  the  weight  of 
silver  in  the  compound.  Thus 

144.34  :  107.88  =  weight  of  AgCl  :  weight  of  Ag. 

This  is  the  gravimetric  method. 

In  the  volumetric  method  the  hydrochloric  acid  would  be 
made  of  such  a  concentration  that  each  cubic  centimeter 
would  contain  a  known  and  definite  weight  of  HC1.  From 
the  number  of  cubic  centimeters  of  acid  used  and  the  con- 
centration of  each  we  can  calculate  the  weight  of  HC1  nec- 
essary to  precipitate  the  silver.  A  proportion  gives  the 
weight  of  silver  precipitated,  this  being  equal  to  107.88  parts 
of  silver  for  each  molecule  of  HC1.  Thus 

36.468  :  107. 88  =  weight  of  HC1  used  :  weight  of  Ag. 

Gravimetric  methods  require  a  considerable  time  for  their 
completion  and  usually  demand  more  manipulative  skill 
than  the  volumetric.  A  very  sensitive  and  somewhat  expen- 
sive balance  is  also  a  necessity.  On  the  other  hand,  the  volu- 
metric methods  can  be  easily  and  rapidly  carried  out  after  the 
standard  solutions  are  prepared  and  the  apparatus  is  neither 
dear  nor  cumbersome.  In  some  cases,  on  account  of  the  im- 
possibility of  deciding  when  the  proper  amount  of  the  reagent 
has  been  used,  volumetric  methods  of  analysis  must  be  aban- 
doned. However,  since  they  are  so  convenient  of  applica- 
tion, they  will  be  the  principal  ones  studied  here. 

The  measuring  apparatus  employed  in  volumetric  chemical 
analysis  is  for  the  determination  of  volume  and  it  is  graduated 
according  to  the  metric  system  (see  table,  page  242)  be- 
cause of  the  great  readiness  by  which  calculations  can  be 
made  thereby.  The  unit  most  in  use  is  the  cubic  centimeter 
(c.c.),  with  its  divisions  and  multiples. 


INTRODUCTION   TO    CHEMICAL   ANALYSIS 


Measuring  cylinders  are  narrow  cylinders  of  glass  with 
a  foot.  On  account  of  the  comparatively  large  surface  of  the 
liquid,  whereby  a  slight  variation  in  the  height  of  the  liquid 
makes  a  considerable  difference  in  its  volume  they  can  only 
be  used  where  accuracy  is  not  essential.  They  will,  however, 
roughly  measure  any  volume  less  than  their  total  capacity. 
Some  are  of  the  same  diameter  throughout  their  length  and 
others  are  fitted  with  glass  stoppers  for  convenience  in  shak- 
ing their  contents. 


FIG.  13. — Apparatus  for  Measuring,  i.  Flasks.  2.  A  pouring  burette 
(Bink's  form).  3.  Two  dropping  burettes  (Mohr's  form)  in  a  holder,  one  with 
a  rubber  tube  and  spring  clamp,  the  other  with  a  glass  stop-cock.  4.  Gradua- 
ted cylinders.  5.  In  front  are  three  pipettes  with  bulbs  and  one  of  the 
cylindrical  form. 

If  accurate  measurements  of  comparatively  large  volumes 
are  necessary  flasks  serve  the  purpose  better.  Accuracy 
is  gained  by  a  narrow  neck  where  they  are  marked  to  indicate 
the  height  to  which  they  should  be  filled.  It  is  advantageous 
to  have  two  marks,  the  lower  one  of  which  shows  where  the 
liquid  should  stand  when  the  flask  contains  the  volume  which 
it  is  designed  to  hold,  for  example,  a  liter,  and  the  upper  (but 
little  above  the  first)  to  which  the  fluid  must  rise  if  the  flask 


VOLUMETRIC   ANALYSIS  151 

is  expected  to  deliver  a  liter  when  it  is  emptied.  That  is, 
allowance  is  made  for  the  few  drops  which  adhere  to  the  in- 
side of  the  vessel.  Measurements  can  by  this  means  be  made 
of  only  one  volume  and  no  fractional  parts  of  it.  Measuring 
flasks  are  used  for  the  preparation  of  the  standard  solutions 
that  are  employed  in  volumetric  analysis.  As  variations  in 
temperature  cause  expansion  and  contraction  in  the  liquids 
the  volume  will  be  correctly  measured  only  at  that  tempera- 
ture at  which  the  flask  was  calibrated.  This  is  usually  at 
15°  C. 

A  pipette  is  a  measure  used  to  deliver  a  definite  and  fixed 
volume  also.  It  is  made  of  a  glass  tube,  either  with  an  expan- 
sion or  bulb  in  the  middle  or  of  the  cylindrical  form.  To  fill  it 
the  point  is  inserted  in  the  liquid,  then  the  air  above  is  ex- 
hausted by  the  mouth  until  the  liquid  rises  above  the  upper 
mark,  when  the  tube  is  removed  from  the  mouth  and  the  top 
quickly  covered  with  the  forefinger.  By  slightly  turning  the 
finger  the  liquid,  which  should  now  stand  above  the  mark,  is 
allowed  to  flow  out  until  the  surface  coincides  with  the  mark. 
The  pipette  is  then  held  over  another  vessel  and  the  liquid  is 
poured  into  this.  Some  pipettes  have  two  graduation  marks, 
one  above  the  bulb  and  the  other  below.  In  this  case  the 
liquid  must  be  allowed  to  escape  only  until  it  reaches  the  level 
of  the  lower  one.  If  there  is  but  a  single  graduation  above 
the  bulb,  it  should  be  emptied,  and  the  last  drop  removed  by 
touching  the  side  of  the  vessel  below.  Pipettes  of  this  shape 
are  very  accurate  since  the  area  of  the  tube  is  small  as  compared 
with  that  of  the  bulb.  Each  one  will  deliver  only  the  volume 
for  which  it  is  calibrated.  They  are  ordinarily  not  used  for 
measuring  over  100  c.c.  The  cylindrical  form  allows  the 
measurement  of  fractions  of  its  whole  volume,  being  gradu- 
ated throughout  the  greater  part  of  its  length.  With  the 
smaller  sizes  a  rubber  bulb  or  nipple  slipped  over  the  upper 
end  enables  the  operator  to  regulate  the  flow  to  dropping 
or  any  desirable  speed. 


152  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

When  measurements  are  to  be  made  of  varying  volumes  of 
liquid  a  burette  should  be  employed.  .  This  is  made  of  a  cylin- 
drical tube,  carefully  graduated  into  cubic  centimeters  and 
their  fractions.  It  differs  from  the  cylindrical  pipette  in  not 
having  its  delivery  regulated  by  the  finger  on  the  upper  end. 
There  are  two  varieties,  the  dropping  and  the  pouring 
burettes.  The  flow  from  the  former  is  usually  regulated  by  a 
valve  or  stop-cock  below.  One  with  a  glass  stop-cock, 
although  the  most  expensive,  is  preferable  since  rubber  affects 
the  strength  of  some  solutions  if  allowed  to  remain  in  contact 
with  them.  (A  little  vaselin  will  prevent  the  stop-cock 
sticking.)  Instead  of  this  in  many  cases  a  short  piece  of 
rubber  tub;ng  can  be  slipped  over  the  end  of  the  burette  and 
the  amount  of  liquid  delivered  can  be  controlled  by  the  com- 
pression of  this  by  a  spring  clamp  or  pinch-cock.  A  solid 
bead  of  glass  slightly  larger  than  the  inside  diameter  will 
also  close  the  tube  when  placed  within.  By  squeezing  this 
between  the  thumb  and  forefinger  the  rubber  is  stretched  and 
the  contents  of  the  burette  slowly  drops  from  the  narrow  jet 
below.  Where  frequent  determinations  are  made  with  the 
same  solution  the  lower  part  of  the  burette  can  be  connected 
with  the  stock  bottle  of  the  same,  thus  permitting  it  to  be 
filled  without  pouring  into  the  top  of  the  tube. 

The  pouring  burettes  vary  somewhat  in  form,  but  are  es- 
sentially a  long,  narrow  tube  graduated  like  the  others  but 
closed  at  the  bottom.  From  these  the  solution  is  poured  out 
of  a  small  jet  at  the  top,  and  the  amount  thus  used  is  indi- 
cated by  the  difference  between  the  height  of  the  liquid  before 
and  after  the  pouring.  By  holding  the  finger  over  the  large 
tube  to  govern  the  admission  of  air  the  flow  can  be  exactly 
regulated.  They  have  no  rubber  parts  to  change  the 
strength  of  the  solutions  used  in  them  and  do  not  cost  as 
much  as  those  with  glass  stop-cocks. 

An  examination  of  the  surface  of  a  liquid  in  a  glass  vessel, 


VOLUMETRIC   ANALYSTS 


153 


particularly  if  this  is  a  narrow  one,  shows  that  the  surface  is 
not  flat  but  concave  or  saucer-shaped.  This  is  called  the 
meniscus.  It  results  from  the  attraction  between  the  glass 
and  liquid  which  causes  a  rising  of  the  latter  at  the  circum- 


FIG.  14.— The  lower  part  of  aMohr's  FIG.  15.— The  lower  part  of  a 
burette  with  spring  clamp,  showing  the  Mohr's  burette  with  glass  stop-cock 
meniscus.  and  Erdmann's  float  in  the  solution. 

ference.  When  seen  from  the  side  it  is  sometimes  a  matter 
of  uncertainty  which  part  should  be  regarded  as  the  top  of  the 
column.  In  measuring  with  a  burette  it  is  a  matter  of  indif- 
ference providing  the  same  part  is  always  used.  With  deeply 


154  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

colored  liquids  only  the  upper  line  can  be  distinguished. 
With  colorless  ones  the  lower  margin  of  the  meniscus  is  most 
distinct,  especially  if  a  white  card  is  held  behind  it.  In  all 
cases  the  eye  of  the  observer  must  be  on  the  same  level  as  the 
meniscus.  The  Erdmann's  float  or  swimmer  may  be  em- 
ployed as  an  aid  in  determining  how  much  the  column  has 
fallen  in  the  burette.  This  is  a  narrow  tube  weighted  so  as  to 
float  vertically  in  the  burette.  A  line  around  the  float  is  the 
measuring  point  instead  of  the  meniscus. 

The  standard  solutions  employed  in  volumetric  analysis 
are  of  two  kinds,  the  normal  and  the  empirical.     A  normal 

/N\ 

solution  fy  1  contains  in  one  liter  a  number  of  grammes  of 

the  reagent  equal  to  its  molecular  weight,  provided  the  mo- 
lecular weight  of  the  reagent  is  equivalent  to  one  atom  of 
hydrogen.  Otherwise  one  liter  contains  such  a  fraction  of 
the  molecular  weight  as  is  the  equivalent  of  one  hydrogen 
atom.  For  example, 

Valence  of  Molecular      Normal  Solu- 


Reagent.        Active  Com- 

Active Com-        Weight. 

tionGrammes 

ponent. 

ponent. 

per  Liter.1 

HC1 

IT 

I 

36.468 

36.468 

NaOH 

OH' 

I 

40  .  008 

40  .  008 

NH4OH 

OH' 

I 

35-05 

35-05 

Na2C03 

CO3" 

II 

106.0 

53-o 

H2SO4 

HJ' 

II 

98.086 

49  .  043 

H2C2O4,2H20 

Hi; 

II 

126.048 

63.024 

H3P04 

Hi" 

III 

98.024 

32-675 

2KMn04 

062 

V 

158-03 

31.606 

Since  there  are  1,000  c.c.  in  a  liter,  each  cubic  centimeter 
of  a  normal  solution  contains  as  many  milligrammes  of  the 
reagent  as  there  are  grammes  in  the  liter.  One  cubic  centi- 
meter of  a  normal  solution  is  the  equivalent  of  one  cubic  centi- 

1  In  making  normal  solutions,  unless  exceptional  accuracy  is  required,  the 
weighings  do  not  need  be  carried  farther  than  two  decimal  places. 

2  KMnO4  acts  as  an  oxidizing  agent,  two  molecules  yielding  5  atoms  of 
oxygen. 


VOLUMETRIC    ANALYSIS  1  55 

meter  of  any  other  normal  solution.     This  is  illustrated  by 
the  following  examples. 

HC+NaOH  =  NaCl+H20. 
36.468  40.008 


.086  80.016 


49.043  40.008 
2KMnO4+5H2C2O4,  2H2O+3H2SO4  = 

10)316.06  630.24 
31.606  63.024 
K2SO4+2MnSO4-f.ioCO2-f8H2O. 

The  equations  represent  the  chemical  actions  taking  place 
between  different  compounds,  the  numbers  below  giving  the 
weight  of  the  one  or  more  molecules  which  enter  into  the  re- 
action. In  all  cases  these  weights  are  seen  to  be  in  the  same 
relation  as  those  in  the  cubic  centimeter  of  the  normal  solu- 
tion, therefore  equal  volumes  of  such  normal  solutions  are 
equivalent  in  their  power. 

In  many  determinations  a  normal  solution  is  too  concen- 
trated for  use  and  one  of  a  fractional  amount  may  be  sub- 

/  N\ 

stituted.     One-tenth  of  the  normal  decinormal,  —  )  or    one 

V  io/ 

/  N  \ 

one-hundredth  of  the  normal  I  centinormal,          )  are  com- 

V  loo/ 

mon.  These  are  prepared  by  placing  in  a  measuring  flask 
one-tenth  or  one  one-hundredth  of  its  volume  of  the  normal 
solution,  then  diluting  to  the  mark.  In  the  same  manner 
others  can  be  made  if  more  convenient. 

An  empirical  standard  solution  is  not  made  up  in  the 
same  manner  as  the  normal,  but  is  represented  by  Fehling's 
solution  for  the  determination  of  glucose  and  some  other 
sugars.  Enough  copper  sulphate  is  taken  so  that  i  c.c.  will 
oxidize  0.005  grm.-of  glucose. 


156  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

The  operation  of  volumetric  analysis  is  called  titration,  and 
the  analyst  speaks  of  titrating  the  substance  under  investi- 
gation. The  process  is  to  place  the  standard  solution  in  a  bu- 
rette, noticing  the  exact  height  of  the  liquid  if  it  does  not 
stand  at  the  zero  mark.  From  this  the  solution  is  allowed  to 
flow  slowly  into  the  liquid  that  is  being  tested,  stirring  con- 
tinually to  cause  immediate  mixing.  When  enough  of  the 
reagent  has  been  added  to  complete  the  reaction  the  amount 
is  carefully  read  from  the  burette  and  the  unknown  quantity 
of  the  other  substance  is  calculated.  In  order  that  a  volu- 
metric method  of  quantitative  analysis  should  be  exact  it  is 
necessary  that  the  end  reaction,  or  point  when  sufficient  of 
the  reagent  is  present,  should  be  plainly  visible  through  some 
change  in  the  appearance.  Sometimes  this  occurs  spon- 
taneously when  the  chemical  change  is  complete,  but  very 
often  a  third  substance  called  an  indicator  must  be  added. 

The  calculation  of  the  results  may  be  made  in  a  number  of 
ways.  It  should  be  remembered  that  although  the  measure- 
ments are  by  volume  the  results  are  in  terms  of  weight  and 
that  therefore  volume  should  be  converted  into  weight  as  soon 
as  possible.  Suppose  that  in  the  determination  of  the  per- 
centage of  silver  in  an  alloy  with  copper  after  dissolving  one 
gramme  of  the  alloy  in  nitric  acid  5.0  c.c.  of  the  normal  hydro- 
chloric acid  should  be  exactly  sufficient  to  precipitate  the 
silver.  Each  c.c.  of  the  normal  solution  containing  0.036468 
grm.  of  HC1,  o.  18234  grm.  of  HC1  would  have  been  thus  used. 
By  the  method  of  proportion  given  on  page  149  we  have 

Molecular  Weight  of  HC1   :  Atomic  Weight  of  Ag  = 
36.468  :  107.88 

Weight  of  HC1   :  Weight  of  Ag 

0.18234  grm.   :  0.5393  grm. 
or  53.93  per  cent,  of  silver  in  the  alloy. 

Or  again,  as  the  equation  on  page  149  shows  that  each  cubic 
centimeter  of  normal  HC1  precipitates  0.10788  grm.  Ag,  the 


VOLUMETRIC    ANALYSIS  1 57 

amount  of  silver  present  can  be  obtained  by  multiplying  this 
by  the  volume  of  the  acid  used.     That  is  (since  normal  solu- 

molecular  weight 

tions   contain  -  of    umvalent    compounds. 

1000 

molecular  weight 

-  of  bivalent  compounds,  etc.,  and  since  one 
2000 

cubic    centimeter    of    this     normal     solution     acts     upon 

molecular  weight 

-  of  other  umvalent  compounds, 

1000 

molecular  weight 

of  other  bivalent  compounds, 
2000 

molecular  weight 

-  of   other   tnvalent   compounds,   etc.)    m 
3000 

volumetric  determinations  multiplication  of  the  number  of 
cubic  centimeters  of  a  normal  solution  necessary  for  a  complete 

molecular  weight   , 
chemical  reaction,  by  -  — of  a  umvalent  compound, 

molecular  weight 

whose  amount  is  being  ascertained  or  by  ~  -  of 

2000 

bivalent  compounds,  etc.,  gives  the  weight  of  the  latter  in  grammes. 


158  INTRODUCTION   TO   CHEMICAL  ANALYSIS 


CHAPTER  II 

ANALYSIS  BY  NEUTRALIZATION 

OF  this  there  are  two  kinds,  acidimetry — the  determination 
of  acids,  or  substances  with  an  acid  reaction — and  alka- 
limetry— the  determination  of  alkalies,  or  substances  having 
an  alkaline  reaction.  A  standard  alkali  or  acid  is  added  from 
the  burette  to  *he  solution  of  unknown  strength  until  the 
reaction  of  the  mixed  solutions  is  neutral.  The  point  is 
shown  by  the  addition  of  an  indicator.  This  is  a  substance 
which  is  of  one  color  in  liquids  of  alkaline  reaction  and 
another  when  the  reaction  is  acid.  A  number  of  such  are 
used  in  acidimetry  and  alkalimetry.  In  some  instances  one  is 
to  be  preferred  and  in  some  another,  so  that  there  can  be  no 
such  thing  as  the  best  indicator  for  all  cases.  Some  from 
their  color,  cannot  be  used  with  liquids  of  a  certain  shade; 
some  have  their  sensitiveness  lessened,  or  their  action  pre- 
vented, by  ammonia,  carbon  dioxid,  etc.;  some  are  affected 
by  mineral  acids,  but  not  by  organic  acids,  sothat  the  selection 
of  the  proper  indicator  is  a  matter  of  importance.  The  prop- 
erties of  some  of  the  most  common  are  given  in  Table  VII. 

The  Preparation  and  Properties  of  Standard  Solutions 

These  can  be  prepared  either  by  standardizing  by  means  of 
other  standard  solutions,  when  it  is  inconvenient  or  im- 
possible to  weigh  the  reagent  which  is  to  be  dissolved,  as  in 
the  case  of  gaseous  reagents  or  those  which  contain  an  indefi- 
nite quantity  of  water.  Or  they  may  be  made,  as  is  most 
frequent,  by  weighing  out  the  proper  amount  of  the  solid 


ANALYSIS  BY  NEUTRALIZATION 


159 


^      . 

"fill 


II 


|31 


C§i 


drtC  13 

ill  I 

II  §  | 

£  o 


l6o  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

reagent,  placing  this  in  a  measuring  flask  (usually  one  holding 
a  liter)  filling  it  to  the  mark  with  distilled  water  of  a  tempera- 
ture of  15°  and  thoroughly  mixing  after  solution  has  occurred. 
When  not  in  use  they  should  be  preserved  in  tightly  closed 
bottles.  Some  slowly  change  for  a  long  time  after  they  have 
been  prepared,  and  others  are  so  unstable  that  special  methods 
of  preservation  must  be  employed  and  their  concentration 
must  be  determined  before  each  new  series  of  analyses.  Too 
great  pains  cannot  be  taken  to  insure  correctness  in  the  stand- 
ard solutions. 

Standard  Solutions  for  Acidimetry  and  Alkalimetry 

Normal  Sodium  Carbonate,  Na2CO^. — Pure  anhydrous 
sodium  carbonate  is  to  be  heated  to  a  low  red  heat  and  after 
cooling,  53.00  grms.  is  dissolved  in  pure  water  and  the  volume 
made  up  to  one  liter.  If  the  pure  salt  is  not  at  hand  85 
grammes  of  sodium  bicarbonate,  which  is  usually  easily  ob- 
tained in  the  pure  state,  will,  by  heating  to  drive  off  the 
water  and  carbon  dioxid,  give  rather  more  than  this  weight 
of  the  pure  carbonate. 

Normal  Sodium  Hydroxid,  NaOH. — This  contains  in 
one  liter  40.008  grammes  of  pure  NaOH.  Since  the  solid 
substance  as  purchased  contains  a  varying  amount  of  mois- 
ture and  consequently  cannot  be  accurately  weighed,  the 
solution  should  be  made  at  first  stronger  than  is  required,  its 
strength  ascertained,  and  should  then  be  diluted  with  the  cal- 
culated volume  of  water.  As  both  the  solid  and  the  solution 
unite  with  carbon  dioxid  care  must  be  taken  to  prevent  its 
access  to  them.  Put  about  45  grammes  of  pure,  dry  sodium 
hydroxid  into  a  liter  flask  and  fill  to  the  mark  with  cold  dis- 
tilled water  from  which  the  carbon  dioxid  has  been  lately  ex- 
pelled by  boiling.  After  dissolving  and  thorough  mixing  re- 
move 10  c.c.  by  a  pipette  to  a  beaker  and  add  a  few  drops  of 
phenolphthalein  solution.  From  a  burette  run  in  normal 


ANALYSIS   BY   NEUTRALIZATION 


161 


hydrochloric  acid  until,  after  stirring,  the  pink  color  has  been 
just  destroyed.     Repeat  this  several  times  and  if  the  deter- 
minations nearly  agree  take  their  average;  calculate  the  con- 
centration of  the  alkaline  solution  and  the  amount  by  which 
it  must  be  diluted  to  make  it  normal.     If,  for  example,  1 1  c.c. 
of    the    normal    acid    is   necessary   to 
neutralize  10  c.c.  of  the  sodium  hydroxid, 
where  if  the  latter  were  normal  only  10 
c.c.    of   acid    would    be    required,    the 
alkali  is  more  concentrated  than  normal. 
Enough  water  must  be  added  to  dilute 
each  10  c.c.   of  the  hydroxid  to  1 1  c.c 
—that   is,   the  one-tenth  of  its  bulk. 
For  example,  if  50  c.c.  of  the  liter  first 
prepared  have  been  used  in  the  prelimi- 
nary testing,  95  c.c.  of  water  should  be 
added  to    the  remainder   (in  a  larger 
flask,    since    this  makes  more  than  a 
liter) .     If  the  determination  was  correct 
10  c.c.  of  this  solution  will  now  exactly 
neutralize    10  c.c.  of  the  normal  acid. 
For  preservation  the  solution  should  be 
kept    in   a  bottle   having  a  two-holed 
rubber    stopper.     Through     one    hole 
passes   a   syphon   tube  with   a  rubber 
tube  and  spring  clamp  on  the  end;  to 
the  other  is  attached  a   U-tube  filled 
with    small    fragments     of   soda-lime. 
The  air  can  pass  through  this,  but  its  carbon  dioxid  is  ab- 
sorbed.    Other  standard  alkaline  solutions,  ammonium,  po- 
tassium  or  barium  hydroxide,  etc.,  can  be  prepared  in  a 
similar  manner. 

Normal    hydrochloric    acid,    HCl,    contains    36.468   grm. 

in  a  liter.     Since  the  acid  as  purchased  is  of  variable  concen- 
11 


FIG.  16. — Bottle  for 
the  preservation  of 
standard  solutions  which 
are  affected  by  the  car- 
bon dioxid  of  the  air. 


1 62  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

tration,  the  normal  acid  must  be  made  by  standardizing  by 
means  of  a  standard  alkaline  solution.  Dilute  about  130  c.c. 
of  the  concentrated  hydrochloric  acid  (specific  gravi  ty  i.i  6)  to 
a  liter,  which  will  make  it  above  normal  concentration.  Mix 
well  and  remove  10  c.c.  to  a  small  beaker,  adding  a  few  drops 
of  methyl-orange  solution.  From  a  burette  allow  a  normal 
sodium  carbonate  solution  to  flow  slowly  in,  stirring  mean- 
while, until  the  color  of  the  liquid  changes  from  a  pink  to  a  pale 
yellow.  Take  the  average  of  several  determinations  and  cal- 
culate the  amount  by  which  it  must  be  diluted  to  make  10  c.c. 
of  one  solution  neutralize  10  c.c.  of  the  other.  Add  this  and 
test  it  to  make  sure  that  it  is  correct.  Thus  if  10  c.c.  of  the 
acid  requires  12.5  c.c.  of  the  normal  carbonate  to  neutralize 
it,  for  each  10  c.c.  of  the  acid  2.5  c.c.  of  water  must  be  added. 
Normal  oxalic  acid,  H^C^O^^H^O,  contains  63.024  grms. 
of  the  crystallized  solid  per  liter.  The  solution  may  be  pre- 
pared by  weighing  out  this  amount  if  the  crystals  are  bright 
in  appearance  and  free  from  moisture.  They  are,  however, 
very  efflorescent,  losing  a  great  part  of  their  water  of  crys- 
tallization when  exposed  to  the  air,  then  becoming  dull  in  ap- 
pearance, and  should  not  be  used  for  making  the  solution  by 
weight.  The  pure  crystals  leave  no  residue  if  they  are  ignited 
on  platinum  foil.  If  such  a  residue  remains  it  is  a  sign  of  an 
impurity.  If  the  pure  crystals  cannot  be  obtained  a  slightly 
greater  weight  than  given  above  may  be  taken  (perhaps  65 
grammes),  and,  after  dissolving  in  water  in  a  liter  flask  and 
diluting  to  a  liter,  the  concentration  may  be  learned  by  titrat- 
ing with  the  normal  sodium  hydroxid,  using  phenolphthalein 
for  the  indicator.  This  solution  will  be  too  concentrated  and 
should  be  reduced  by  adding  the  calculated  amount  of  water 
as  in  the  preparation  of  normal  hydrochloric  acid.  The  solu- 
tion decomposes  in  direct  sunlight.  Very  dilute  solutions  do 
the  same  in  weak  light  so  that  if  decinormal  or  centinormal 
solutions  are  needed  the  normal  solution  should  only  be 
diluted  with  water  shortly  before  it  is  to  be  made  use  of. 


ANALYSIS  BY   NEUTRALIZATION  163 

Other  acids,  sulphuric,  tartaric,  etc.,  can  be  standardized 
by  a  similar  course  of  procedure. 

Acidimetry 

For  the  standard  solution  normal  sodium  carbonate  or 
sodium  hydroxid  may  be  used.  Place  this  in  the  burette, 
making  a  note  of  its  height  if  it  is  not  at  the  zero  mark,  and 
with  a  pipette  carefully  measure  a  definite  volume,  perhaps  10 
c.c.  of  the  liquid  of  which  the  acidity  is  to  be  determined, 
transferring  it  to  a  small  beaker.  To  this  add  a  few  drops  of 
the  indicator,  selected  in  accordance  with  Table  VII.  For 
the  titration  let  the  standard  alkali  flow  slowly  into  the  acid, 
stirring  meanwhile  in  order  to  thoroughly  mix  the  two  liquids. 
At  first  the  color,  produced  when  the  alkali  meets  the  indi- 
cator, vanishes  immediately.  When  it  disappears  only  slowly 
the  normal  solution  should  be  added  by  drops  so  as  to  avoid 
an  excess,  stopping  when  the  change  in  color  of  the  indicator 
is  persistent  throughout  the  liquid.  The  amount  drawn  from 
the  burette  should  be  ascertained  and  noted  and  the  titration 
repeated.  It  is  best  to  take  the  average  of  several  nearly 
concordant  results  than  to  depend  upon  one.  Knowing  the 
weight  of  each  acid  which  will  be  neutralized  by  one  cubic 
centimeter,  the  weight  present  in  the  volume  used  can  be  cal- 
culated. The  final  results  can  be  expressed  as  grammes  per 
liter  or  in  percentage  by  weight.  In  a  liter  of  water  there  are 
1,000  grammes  (c.c.)  and  consequently  if  the  liquid  being 
tested  is  of  the  same  specific  gravity  as  water  the  percentage 
by  weight  (that  is,  parts  in  a  hundred)  can  be  obtained  from 
the  grammes  per  liter  by  dividing  by  ten.  When  the  specific 
gravity  is  greater  or  less  than  one  the  result  so  obtained  must 
be  divided  by  the  specific  gravity  to  find  percentage.  It  is  to 
be  understood,  of  course,  that  acids  combined  in  neutral  salts 
cannot  be  determined  in  this  manner. 


164  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

Practical  Exercises  in  Acidimetry 

1 .  Determine  the  concentration  of  dilute  sulphuric  and  nitric 
acids  furnished  by  the  instructors,  reporting  results.     Use 
normal  sodium  carbonate,  with  methyl  orange  for  an  indicator.1 

2.  From  concentrated  hydrochloric  acid  prepare  a  normal 
solution  according  to  the  directions  on  page  161. 

3.  After  the  preparation  of  normal  sodium  hydroxid  (page 
1 60)  standardize  a  normal  solution  of  oxalic  acid  (page  162). 

Other  chemical  compounds  which  have  an  acid  reaction  can 
be  determined  quantitatively  by  similar  methods  after  the 
choice  of  a  suitable  indicator.  Among  those  with  which 
sodium  hydroxid  and  phenolphthalein  can  be  used  are  acetic 
acid  (vinegar),  lactic  acid  (sour  milk)  or  acidity  of  gastric  con- 

N 
tents  (with —  NaOH),  aromatic  sulphuric  acid,  hydrobromic 

acid,  hypophosphorous  acid,  HPH202  (univalent,  the  equa- 
tion being 

NaOH  +  HPH2O2  =  NaPH2O2  +  H2O). 

In  all  cases  the  basicity  of  the  acid — that  is,  the  number  of 
hydrogen  atoms  which  can  be  ionized,  must  be  considered,  as 
this  determines  the  amount  of  alkali  necessary  to  neutralize 
the  molecule. 

Alkalimetry 

In  this  the  operations  of  acidimetry  are  reversed,  the  stand- 
ard acid  solution  being  placed  in  the  burette  and  the  alkaline 
solution  below.  Otherwise  the  method  is  as  before.  In  ad- 
dition to  its  use  in  determining  the  amount  present  of  a  sub- 
stance with  an  alkaline  reaction  it  can  be  used  for  the  same 
purpose  with  another  class  of  compounds — that  is,  the  salts  of 

1  In  the  practical  exercises  a  single  titration  ought  not  to  be  depended  upon 
but  the  average  of  two  or  three  closely  agreeing  ones  should  be  taken. 


ANALYSIS  BY  NEUTRALIZATION  165 

organic  acids  which  have  as  a  base  one  of  the  alkali  metals. 
These  when  ignited  are  changed  into  carbonates  of  those 
metals.  The  carbonates  are  alkaline  in  reaction  although  the 
salts  from  which  they  are  derived  may  be  acid  or  neutral. 
One  molecule  of  the  carbonate  requires  for  its  formation  a 
sufficient  number  of  molecules  of  the  organic  compound  to 
furnish  two  atoms  of  a  monovalent  metal.  Thus  one  mole- 
cule of  a  normal  tartrate  gives  one  molecule  of  carbonate, 


Two  molecules  of  an  acetate  or  acid  tartrate     form  the  same 
amount. 


Two  molecules  of  a  normal  citrate  form  three  of  a  carbonate. 


It  will  not  be  necessary  to  burn  off  all  the  carbon  to  form  the 
carbonate  —  that  is,  to  heat  long  enough  to  leave  a  perfectly 
white  residue.  If  the  organic  salt  is  at  first  in  solution  a 
measured  amount  must  be  evaporated  to  dryness,  then  the 
solid  must  be  heated  in  a  platinum  or  porcelain  crucible, 
gently  at  first,  to  avoid  loss,  then  until  the  mass  is  completely 
carbonized,  which  will  require  a  red  heat.  If  the  compound 
is  a  solid,  a  weighed  quantity  (one  to  two  grammes)  should  be 
treated  in  the  same  way.  After  cooling  the  carbonate  will 
dissolve  easily  in  hot  water.  It  should  be  filtered  and  the 
carbon  washed  with  small  amounts  of  hot  water  until  it  is  no 
longer  alkaline.  The  weight  of  carbonate  can  be  determined 
by  titration  and,  knowing  the  number  of  molecules  of  the  or- 
ganic^salt  from  which  one  of  the  carbonate  is  derived  together 
with  their  relative  molecular  weights,  the  weight  of  the 
original  salt  is  readily  calculated. 


1 66  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

Practical  Exercises  in  Alkalimetry 

1 .  Using  normal  hydrochloric  acid  for  the  standard,  deter- 
mine the  concentration  of  an  unknown  solution  of  ammonium 
hydroxid  with  cochineal  or  rosolic  acid  as  indicator ;  of  borax 
with  lacmoid  as  an  indicator. 

2.  Prepare  a  normal  solution  of  ammonium  hydroxid  by 
the  method  outlined  on  page  161. 

3.  Ignite  thoroughly  in  a  platinum  or  porcelain  capsule 
a  weighed  amount  of  acid  potassium  tartrate,  extract  with 
hot  water,  titrate  the  nitrate,  using  normal  hydrochloric  acid 
and  methyl  orange,  and  from  calculated  result  form  an  opin- 
ion as  to  the  purity  of  the  tartrate. 

4.  Evaporate  to  dryness  in  a  porcelain  capsule  10  c.c.  of  a 
concentrated  solution  of  sodium  acetate,  ignite,  titrate  the 
soluble  part  and  calculate  the  concentration  of  the  solution. 

Among  the  other  compounds  of  which  the  amounts  can  be 
found  by  alkalimetry  are  the  hydroxids  and  carbonates  of  the 
alkali  metals,  the  hydroxids  of  the  alkaline  earths,  the  tar- 
trates  of  sodium  and  potassium,  the  acetate  of  potassium,  the 
citrates  of  potassium  and  lithium,  the  benzoates  of  lithium 
and  sodium,  the  salicylates  of  lithium  and  sodium. 


ANALYSIS  BY   OXIDATION   AND  BY  REDUCTION          167 


CHAPTER  III 

ANALYSIS  BY   OXIDATION    (OR   OXIDIMETRY),   AND  BY 
REDUCTION 

THE  method  is  based  upon  the  fact  that  certain  compounds 
have  the  power  of  oxidizing  definite  amounts  of  other  com- 
pounds when  in  solution.  In  some  cases  the  end  of  the 
reaction  is  indicated  by  a  change  in  the  color  of  the  standard; 
in  others  indicators  are  added.  The  results  obtained  are  very 
accurate.  The  most  common  oxidizing  agents  used  for  the 
standard  solutions  are  potassium  permanganate,  KMnO4; 
potassium  dichromate,  K^C^O?,  and  iodin. 

The  Preparation  and  Properties  of  Standard  Oxidizing 
Solutions 

Potassium  permanganate,  KMnO^  can  be  obtained  pure  in 
the  crystalline  form,  though  it  is  so  often  impure  that  the  solu- 
tion should  be  tested  after  making  by  titrating  against  a  stand- 
ard. The  crystals  dissolve  to  an  intensely  reddish-purple  solu- 
tion. The  molecular  weight  of  the  substance  is  158.03.  As 
stated  before  (page  154)  when  it  acts  as  an  oxidizing  agent  two 
molecules  yield  five  atoms  of  oxygen.  The  combined  valences 
of  the  five  atoms  of  oxygen  given  by  this  double  molecule  is  ten. 
Therefore  a  normal  solution  would  contain  one-tenth  the 
double  molecular  weight  of  the  salt  in  a  liter  or  3 1 .606  grammes. 
The  decinormal  solution  containing  31.606  is  more  commonly 
employed.  This  solution  is  likely  to  undergo  slight  decom- 
position on  standing,  although  it  will  maintain  its  strength  for 
several  weeks.  Unless  freshly  prepared  it  should  be  standard- 
ized before  the  solution  is  used,  in  the  same  manner  as  in  the 


1 68  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

original  preparation.  Organic  matter  produces  changes  in 
the  solution,  which  should  therefore  not  be  brought  into  con- 
tact with  rubber.  A  pouring  burette  or  one  with  a  glass  stop- 
cock is  suitable  for  its  measurement.  When  the  permanganate 
gives  up  its  oxygen,  providing  a  free  mineral  acid  is  present,  the 
purple  color  disappears  as  a  result  of  the  formation  of  the  color- 
less manganese  salt  of  the  acid.  That  is,  the  manganese, 
which  was  a  part  of  the  anion,  becomes  the  cation.  This 
decolorization  indicates  the  completion  of  the  reaction.  An 
indicator  is  therefore  unnecessary.  If  the  crystals  are  known 
to  be  pure  the  weighed  amount  can  be  dissolved  in  a  liter  of 
water.  In  case  of  doubt  as  to  their  purity  about  3.5  grammes 
should  be  so  dissolved,  the  concentration  of  the  solution  as- 
certained by  one  of  the  following  methods  and  the  proper 
volume  of  water  added  to  reduce  it  to  the  desired  concentration. 

The  substances  most  commonly  used  to  ascertain  the  con- 
centration of  a  permanganate  solution  are  oxalic  acid,  ammo- 
nium ferrous  sulphate,  and  metallic  iron. 

A  decinormal  solution  of  oxalic  acid  can  be  taken  as  the 
standard.  The  reaction  is 

2KMnO4  +  5H2C204  +  3H2SO4 
=  K2S04  +  2MnS04  +  ioC02  +  8H2O. 

The  10  c.c.  of  the  oxalic  acid  is  measured  with  a  pipette  and 
after  acidifying  with  dilute  sulphuric  acid,  is  placed  in  a  flask 
or  beaker  and  the  mixture  warmed  to  about  60°.  Then  the 
permanganate  solution  is  allowed  to  run  in.  At  first  the  color 
disappears  slowly  but  afterward  more  rapidly.  If  it  turns 
brown  the  amount  of  sulphuric  acid  is  insufficient.  The  per- 
manganate should  be  added  cautiously  to  avoid  an  excess, 
stopping  when  the  pink  color  is  permanent.  From  the 
average  of  several  determinations  calculate  the  concentration 
of  the  permanganate  solution,  remembering  that  if  it  were 
decinormal  exactly  10  c.c.  would  be  reduced  by  the  oxalic  acid. 


ANALYSIS  BY   OXIDATION   AND  BY   REDUCTION          169 

Find  the  volume  of  water  which  will,  if  added,  produce  the 
decinormal  solution,  as  in  the  preparation  of  standard  solutions 
of  sodium  hydroxid  (page  160)  or  hydrochloric  acid  (page  161). 
After  dilution  test  again  to  see  if  the  result  is  correct.  It  is 
well  in  this  second  test  to  make  use  of  another  standard  solu- 
tion such  as  one  of  decinormal  ferrous  ammonium  sulphate. 

Ferrous  ammonium  sulphate,  (NH^fe^SO^^  6#20,  can 
be  obtained  pure  in  the  form  of  greenish  crystals.  For  this 
purpose  they  must  not  have  lost  any  of  their  water  by  efflores- 
cence or  be  at  all  brown  in  color,  which  indicates  that  the  iron 
is  changing  to  the  ferric  form.  A  decinormal  solution  con- 
taining 38.41  grams  to  the  liter  should  be  used  with  sulphuric 
acid  in  the  same  manner  as  in  the  titration  by  oxalic  acid. 
The  reaction  is 


Here  as  before  one  cubic  centimeter  of  the  iron  solution 
should  decolorize  exactly  one  of  the  decinormal  potassium 
permanganate. 

If  metallic  iron  is  the  standard  about  o.i  grm.  of  the  purest 
piano  wire,  accurately  weighed,  should  be  dissolved  in  a  flask 
by  means  of  dilute  sulphuric  acid.  Access  of  air  can  be  pre- 
vented by  a  Bunsen  valve  made  of  a  short  piece  of  slit  rubber 
tubing  slipped  over  the  exit  tube  and  closed  at  the  upper  end, 
which  allows  the  hydrogen  to  escape  (Fig.  17).  When  the 
wire  has  dissolved,  the  solution  can  be  diluted  with  recently 
boiled  water  and  immediately  titrated  with  the  permanganate. 
The  reaction  is  similar  to  that  with  the  ammonium  ferrous 
sulphate  because  the  iron  in  dissolving  forms  ferrous  sul- 
phate, FeSO4.  Since  all  iron  contains  carbon,  allowance 
must  be  made  for  this.  It  it  has  not  been  determined  in  the 
sample  which  was  used,  the  wire  may  be  estimated  at  99.6  per 
cent.  pure. 


170 


INTRODUCTION   TO   CHEMICAL  ANALYSIS 


Of  iodin  the  decinormal  solution  is  used.  This  contains 
1 2 .692  grms.  in  the  liter.  It  can  be  made  by  dissolving  about 
1 8  grms.  of  pure  potassium  iodid  in  200-300  c.c.  of  water  in  a 
liter  flask,  adding  12.692  grms.  of  chemi- 
cally pure  iodin  and  when  this  has  dis- 
solved, filling  to  the  mark  with  water.  If 
pure  iodin  cannot  be  obtained  somewhat 
more  than  this  weight  may  be  dissolved 
and  the  concentration  of  the  solution  be 
ascertained  by  titrating  with  a  standard 
solution  of  sodium  thiosulphate  (hypo- 
sulphite) in  the  manner  described  below. 
The  iodin  solution  should  be  kept  in  a 
cool,  dark  place,  but  even  then  it  does 
not  maintain  its  concentration,  so  that 
unless  it  has  been  recently  prepared  it 
should  be  standardized  before  using. 

The  oxidizing  action  of  the  iodin  is  an 
indirect  one.  It  unites  with  the  hydrogen 
of  the  water  present,  leaving  the  oxygen 
free  to  combine  with  oxidizable  sub- 
stances. This  is  illustrated  by  its  action 


FIG.  17.  — Flask 
fitted  with  Bunsen 
valve  which  allows 
the  gas  to  escape  from 


within  but  prevents    on  sociium  sulphite  or  arsenous  oxid. 

the  access  of  air. 

Na2S03+I2+H20  =  2HI+Na2S04. 


Although  solutions  of  free  iodin  have  a  color  and  most 
solutions  of  its  compounds  do  not,  the  distinction  is  not  great 
enough  to  accurately  mark  the  end  reaction.  Therefore  a 
few  drops  of  a  boiled  starch  solution  are  added  as  an  indica- 
tor. This  gives  a  deep  blue  color  as  long  as  any  free  iodin  is 
present  and  is  colorless  when  the  iodin  is  in  combination  with 
other  elements. 

Decinormal  potassium  dichr  ornate,  K^Cr^O?,  which  is  often 


ANALYSIS  BY   OXIDATION   AND  BY  REDUCTION          171 

used  in  oxidimetry,  contains  4.907  grms.  in  a  liter.  The  crys- 
tallized salt  can  be  obtained  in  a  pure  state  and  the  solution 
may  be  made  by  drying  the  salt  at  100°,  then  dissolving  this 
weight  in  water  and  diluting  to  a  liter.  The  solution  is  much 
more  permanent  than  that  of  potassium  permanganate  or  iodin 
and  is  not  affected  by  contact  with  rubber.  An  indicator  must 
be  used  with  it  in  the  volumetric  tests  of  ferrous  compounds 
and  this  lessens  its  convenience.  When  the  dichromate  is 
used  as  an  oxidizing  agent  it  loses  three  atoms  of  oxygen.  It 
is  for  this  reason  that  its  decinormal  solution  contains  one-six- 
tieth of  its  molecular  weight  in  grammes  per  liter  (page  1  54)  . 
The  sulphuric  acid  combines  with  the  potassium  and  chromium 
forming  the  sulphates  of  these  metals.  Thus 


The  principal  use  of  the  dichromate  solution  in  volumetric 
analysis  is  in  the  quantitative  determination  of  iron  and  its 
ferrous  compounds.  If  a  ferrous  compound  is  present  with 
an  acid  the  oxygen  unites  with  the  hydrogen  ion  of  the  latter, 
and  the  anion  combines  with  iron,  changing  it  to  the  ferric 
state. 


Analysis  of  Reduction 

This  is  the  opposite  of  oxidimetry.  Soluble  reducing 
agents,  or  deoxidizers  will  remove  oxygen  from  many  of  its 
compounds  when  in  solution  and  if  the  end  of  this  reducing 
action  is  definitely  marked  they  may  often  be  employed  in  the 
preparation  of  standard  volumetric  solutions.  Oxalic  acid 
and  sodium  thiosulphate  are  very  commonly  used.  The 
former  has  already  been  discussed. 

Sodium  thiosulphate,  -/Va2S203,5#20,,  contains  in  a  liter  of 
the  decinormal  solution  24.822  grms.  of  the  crystallized  salt. 
The  solution  can  be  prepared  by  dissolving  in  water  this 


172  INTRODUCTION    TO    CHEMICAL    ANALYSIS 

weight  of  the  pure  crystals  which  have  been  dried  by  pressing 
in  blotting-paper  after  pulverizing,  then  diluting  to  a  liter. 
The  solution  should  be  kept  in  the  dark.  It  even  then  de- 
composes slowly,  however,  so  that  old  solutions  cannot  be 
depended  upon.  Sodium  thiosulphate  is  principally  used  in 
volumetric  analysis  in  the  determination  of  free  iodin.  This 
includes  the  determination  of  other  substances  like  bromin 
and  chlorin,  one  atom  of  either  of  which  sets  free  one  atom 
of  iodin  from  an  iodid.  If,  therefore,  to  a  solution  of  free 
bromin  or  chlorin  a  little  potassium  iodid  is  added,  the  bro- 
min or  chlorin  frees  the  iodin.  The  amount  of  iodin  thus 
liberated  indicates  that  of  bromin  or  chlorin. 

Practical  Exercises  in  Analysis  by  Oxidation  and  Reduction 

1 .  From  the  normal  solution  of  oxalic  acid  previously  pre- 
pared make  a  decinormal  solution  by  diluting  one  volume 
with  nine  of  water.     By  the  aid  of  this  standardize  a  decinor- 
mal  solution  of   potassium  permanganate  by  the  method 
described  on  page  168. 

2.  With  the  decinormal  permanganate  makes  a  determina- 
tion of  the  amount  of  iron  in  a  solution  of  ferrous  sulphate, 
calculating  the  weight  of  Fe  and  FeSO4  present. 

3.  By  the  aid  of  a  decinormal  solution  of  potassium  di- 
chromate  (made  by  the  directions  given)  make  a  titration  of 
the  same  ferrous  solution.     Determine  the  end  of  the  reaction 
by  removing  a  small  drop  of  the  solution  on  the  end  of  a  glass 
rod  and  with  this  stirring  a  drop  of  a  dilute,  freshly  prepared 
solution  of  potassium  ferricyanid  as  an  indicator.     As  long 
as  there  remains  any  of  the  unoxidized  iron  a  blue  color  will 
result.     When  sufficient  dichromate  is  present  only  a  brown- 
ish-yellow appears.     The  tests  are  most  conveniently  made 
by  placing  a  number  of  drops  of  the  indicator  on  a  porcelain 
plate  and  touching  these  with  the  stirring  rod  after  the  addi- 
tion of  each  portion  of  the  dichromate.     The  results  obtained 


ANALYSIS   BY    OXIDATION   AND   BY   REDUCTION  173 

by  this  method  should  agree  with  those  by  the  permanganate. 
They  thus  serve  to  confirm  the  correctness  of  the  dichromate 
solution  as  well  as  of  the  accuracy  of  the  determination  of  the 
amount  of  iron. 

4.  Use  the  decinormal  permanganate  to  determine  the  con- 
centration of  a  solution  of  hydrogen  peroxid  (dioxid).     The 
reaction    is    5H2O2+2KMnO4  +  3H2SO4  =  K2SO4+2MnSO4 
-f  8H2O+5C>2.     Measure  into  a  beaker  one  cubic  centime- 
ter of  the  peroxid  by  means  of  a  pipette  and  dilute  with  10 
to  20  times  its  volume  of  water  acidified  with  H2SO4.     From 
a  burette  add  the  permanganate  slowly  until  there  is  a  per- 
manent pink  color.     (In  the  U.  S.  P.  process  10  c.c.  of  the 
peroxid  are  diluted  to  100  c.c.,  and  16.9  c.c.  of  this  is  titrated. 
The  Pharmacopoeia  requires  it  to  decolorize  30  c.c.  of  the 
decinormal  permanganate.     This  corresponds  to  a  3  per  cent, 
solution  by  weight.)     Calculate  the  percentage  strength  of 
the  hydrogen  peroxid  solution  by  weight. 

Instead  of  being  expressed  by  weight  the  concentration  of 
the  hydrogen  peroxid  is  often  referred  to  the  volume  of  oxy- 
gen which  it  will  evolve  when  decomposed  by  heating.  One 
atom  is  thus  set  free  from  each  molecule,  or  one-half  the 
amount  that  is  given  off  when  it  is  acted  upon  by  potassium 
permanganate,  as  represented  by  the  above  equation.  For 
each  two  molecules  of  the  permanganate  which  are  decolorized, 
therefore,  there  are  present  five  atoms  of  active  oxygen  in  the 
peroxid.  Consequently  i  c.c.  of  the  decinormal  permanga- 
nate corresponds  to  0.0008  grm.  of  such  active  oxygen. 
From  the  results  obtained  in  the  above  determination  of  the 
concentration  of  hydrogen  peroxid  by  weight  calculate  its 
oxygen  volume  of  active  oxygen,  using  0.00143  grm.  as  the 
weight  of  one  cubic  centimeter  of  oxygen. 

5.  Iron  is  determined  quantitatively  by  permanganate  or 
dichromate  solutions,  but  only  when  it  is  in  the  ferrous  form. 
Hence  ferric  compounds  must  be  reduced  to  ferrous  before 


174  INTRODUCTION   TO   CHEMICAL  ANALYSIS 

they  are  titrated.  This  can  be  effected  by  a  number  of  reduc- 
ing agents.  Place  in  a  flask  fitted  with  a  valve  as  described 
above  (page  170)  10  c.c.  of  a  ferric  solution,  acidify  with  sul- 
phuric acid  and  add  a  few  small  fragments  of  granulated  zinc 
which  is  free  from  iron  or  in  which  the  amount  of  iron  is 
known  (using  than  a  definite  weight  of  zinc).  Let  it  dissolve 
completely  when,  if  the  iron  is  reduced,  the  liquid  will  be 
colorless  with  no  yellow  tint.  Then  titrate  immediately  with 
the  decinormal  permanganate  solution  and,  after  the  amount 
of  iron  has  been  found,  calculate  the  weight  and  percentage 
of  the  ferric  compound  in  the  original  solution. 

6.  With  the  aid  of  a  decinormal  solution  of  sodium  thio- 
sulphate  determine  the  concentration  of  a  solution  of  iodin, 
as  follows :     Into  a  measured  volume  of  the  iodin  solution  run, 
from  a  burette,  the  standard  thiosulphate  until  the  brown 
color  has  almost  disappeared.     Then  add  a  few  drops  of  a 
starch  solution  and  continue  the  titration  until  the  blue  is 
just  destroyed,  leaving  the  liquid  colorless.     The  reaction  is 
then  completed,  the  thiosulphate  being  converted  into  a 
sodium  tetrathionate, 

2Na2S2O3+I2  =  2NaI+Na2S4O6. 

7.  Chlorin  or  bromin  when  brought  into  contact  with  potas- 
sium iodid  liberates  an  equal  number  of  atoms.     In  conse- 
quence of  this  action  the  concentration  of  chlorin  or  bromin 
water  is  easily  found.     To  10  c.c.  of  the  solution  add  about 
half  a  gramme  of  potassium  iodid  in  crystals  or  solution  and 
titrate  with  thiosulphate,  using  starch  as  an  indicator,  as  in 
the  last  operation. 

8.  In  the  same  manner  determine  the  amount  of  available 
chlorin  in  calcium  hypochlorite,  CaOCl2. 

CaOCl2+2HCl  =  CaCl2+H20+Cl2. 
Acidify  a  known  quantity  with  hydrochloric  acid  after  the 


ANALYSIS  BY   OXIDATION   AND  BY   REDUCTION          175 

addition  of  potassium  iodid  and  starch,  and  titrate  with  the 
thiosulphate  as  in  the  preceding  exercises. 

The  presence  of  chlorates  lessens  the  accuracy  of  the  last 
determinations. 

9.  With  starch  as  an  indicator  compare  the  concentration 
of  the  decinormal  thiosulphate  with  the  decinormal  iodin 
solution  to  prove  that  both  are  correct. 

10.  When  a  ferric  salt  is  warmed  with  potassium  iodid  it  is 
changed  to  the  ferrous  state,  one  atom  of  iodin  being  set  free 
for  each  atom  of  iron. 

FeCl3+KI  =  FeCl2+KCH-I. 

The  quantitative  method  of  determining  ferric  compounds 
based  upon  this  reaction  is  carried  out  in  the  following  manner : 

To  10  c.c.  of  the  ferric  solution  add  about  a  gramme  of  po- 
tassium iodid  and  2  c.c.  of  hydrochloric  acid.  By  the  above 
reaction  the  iodin  is  slowly  set  free.  The  mixture  should  be 
placed  in  a  100  c.c.  glass-stoppered  bottle  and  the  whole 
warmed  in  water  two  hours  at  40°.  The  temperature  should 
not  be  allowed  to  exceed  this  nor  the  stopper  be  removed  be- 
cause of  danger  of  loss  of  iodin  through  volatilization.  Cool, 
and  after  the  addition  of  a  few  drops  of  starch  solution,  titrate 
with  decinormal  sodium  thiosulphate.  Calculate  the  amount 
of  the  ferric  salt,  reckoning  one  atom  of  iron  for  each  one  of 
free  iodin. 

11.  Prepare  a  decinormal  solution  of  iodin  by  weighing 
(page  170),  or  take  a  larger  amount  than  is  necessary,  dissolve 
in  the  same  manner,  determine  its  concentration  by  titration, 
and  dilute  to  the  standard. 

1 2 .  With  the  standard  iodin  solution  make  a  determination 
of  a  solution  of  arsenous  acid,  HsAsOs,  or  an  arsenite,  like 
potassium  arsenite,  KsAsOs.     In  the  case  of  the  arsenous  acid 
pure  sodium  or  potassium  bicarbonate  must  be  present  to 
neutralize  the  hydriodic  acid  formed  in  the  titration.     Five  to 


176  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

ten  times  as  much  of  this  should  be  used  as  the  estimated 
weight  of  the  arsenous  acid.  The  titration  is  carried  out  as 
before  until  the  starch  indicator  is  colored  a  faint  blue.  A 
gentle  heat  may  be  used  to  aid  in  the  solution  of  the  bicarbon- 
ate, but  it  cannot  be  heated  high  enough  to  decompose  with 
bicarbonate  as  the  carbonate  thus  produced  would  interfere 
with  the  action  of  the  indicator.  If  a  free  alkali  other  than  a 
bicarbonate  is  present  it  must  be  neutralized  by  hydrochloric 
acid  and  any  acids  must  be  neutralized  with  a  bicarbonate 
before  titration.  The  reaction  between  the  iodin  and  arsen- 
ite  is  similar  to  that  with  the  arsenous  oxid  already  referred  to. 


ANALYSIS  BY  PRECIPITATION  177 

CHAPTER  IV 

ANALYSIS  BY  PRECIPITATION 

HERE  the  standard  solution  converts  the  compound  of 
which  the  amount  is  sought,  or  some  constituent  of  it,  into  an 
insoluble  form,  thus  producing  a  precipitate.  Knowing  the 
amount  of  the  standard  necessary  to  effect  this  result  the 
weight  of  the  precipitate  or  of  the  substance  from  which  it  is 
derived  can  be  calculated.  As  the  point  where  precipitation 
is  complete  is,  in  almost  all  cases,  indistinct,  indicators  are 
usually  required  in  this  class  of  analysis.  These  are  such 
compounds  as  will  not  be  acted  upon  by  the  standard  until 
the  compound  under  investigation  has  been  completely  trans- 
formed and  which  will  then  react  with  the  excess  of  the  stand- 
ard solution  producing  a  color  or  some  other  visible  change. 

Precipitation  is  also  often  used  in  combination  with  other 
methods  of  volumetric  determination.  Thus  soluble  com- 
pounds of  barium,  strontium,  and  calcium  can  be  precipitated 
as  neutral  carbonates  by  ammonia  and  ammonium  carbonate. 
If  these  precipitates  are  washed  and  suspended  in  water  they 
can  be  titrated  by  solutions  of  the  normal  acids,  their  amount 
being  calculated  from  the  volume  of  acid  necessary  to  produce 
an  acid  reaction.  The  equations  representing  the  chemical 
change  occurring  during  their  solution  may  be  represented  by 
the  following: 


The  Preparation  and  Properties   of   Standard   Solutions 
Used  in  Analysis  by  Precipitation 

Decinormal  Silver  Nitrate,  AgNO^.  —  As  the  crystals  can 

usually  be  obtained  in  the  pure  state,  this  can  be  made  by  dis- 
12 


178  INTRODUCTION   TO    CHEMICAL  ANALYSIS 

solving  16,989  grms.  in  sufficient  water  to  make  the  volume  of 
the  solution  one  liter.  Or  a  somewhat  greater  weight  may  be 
dissolved,  if  the  purity  is  doubtful,  and  the  concentration  as- 
certained by  titration  against  decinormal  sodium  chlorid. 
Instead  of  the  sodium  chlorid  decinormal  hydrochloric  acid 
can  be  used.  In  the  latter  case  after  the  amount  to  be  used 
has  been  accurately  measured  by  a  pipette,  it  must  be  care- 
fully neutralized  with  sodium  carbonate  before  the  titration, 
avoiding  an  excess  of  the  carbonate.  The  value  of  the  silver 
nitrate  in  analysis  by  precipitation  is  that  it  forms  insoluble 
compounds  with  the  chlorids,  bromids,  iodids,  and  cyanids. 
The  indicator  is  generally  normal  (yellow)  potassium  chro- 
mate.  This  forms  an  insoluble,  dark  red,  silver  chromate 
with  the  silver  nitrate,  but  not  until  the  above-mentioned 
compounds  have  been  precipitated  if  they  are  present  in  the 
solution.  The  end  reaction  is  most  clearly  seen  if  instead  of 
daylight  a  yellow  light,  like  that  of  illuminating  gas,  is  used, 
the  titration  being  conducted  in  a  rather  dark  place. 

The  silver  nitrate  solution  is  decomposed  by  the  action  of 
light  and  by  organic  matter.  It  should  therefore  be  pre- 
served in  an  amber-colored  bottle  or  in  a  dark  place.  It 
should  be  protected  from  dust  and  not  be  used  in  burettes 
which  have  rubber-tubing  connections. 

Decinormal  Potassium  Sulphocyanate,  KSCN,  contains  in  a 
liter  9.718  grms.  of  the  salt.  It  cannot  well  be  prepared  by 
weighing  the  solid,  since  the  latter  is  deliquescent.  About 
10  grammes  of  this  should  be  dissolved  in  a  liter  of  water  and 
10  c.c.  of  decinormal  silver  nitrate  titrated  with  the  solution 
after  acidifying  with  5  c.c.  of  dilute  nitric  .acid.  The  reaction 
is  shown  by  the  following  equation: 

AgN03+KSCN  =  AgSCN+KN03. 

About  ten  drops  of  a  solution  of  ammonium  ferric  sul- 
phate (iron  alum)  is  to  be  added  for  an  indicator.  The 


ANALYSIS   BY   PRECIPITATION  179 

silver  sulphocyanate  is  precipitated  first,  giving  the  liquid 
a  milky  appearance.  When  all  the  silver  has  been  con- 
verted into  this  compound  the  sulphocyanate  acts  on  the  in- 
dicator producing  red  ferric  sulphocyanate  which  indicates 
the  end  reaction.  With  a  solution  prepared  in  this  way  less 
than  10  c.c.  of  the  sulphocyanate  should  at  first  precipitate 
10  c.c.  of  the  decinormal  silver  nitrate.  The  amount  of  water 
which  must  be  added  to  dilute  it  to  correspond  to  the  silver 
solution  can  be  calculated  as  in  the  preparation  of  standard 
solutions  of  sodium  hydroxid  and  hydrochloric  acid.  When 
this  has  been  added  it  should  be  tested  again  in  the  same  man- 
ner to  ascertain  its  correctness.  This  sulphocyanate  solution 
may  be  used  for  the  estimation  of  silver,  even  in  the  presence, 
of  many  other  dissolved  metals,  since  most  of  these  are  un- 
affected by  the  reagent.  It  can  also  be  employed  in  connec- 
tion with  decinormal  silver  nitrate  to  determine  the  amount 
of  any  substance  which  is  completely  precipitated  by  the  lat- 
ter compound.  This  includes  all  those  mentioned  as  capable 
of  being  determined  by  standard  silver  nitrate.  In  this  case 
the  method  is  that  of  residual  titration.  It  consists  in  adding 
to  the  solution  a  measured  quantity  of  the  decinormal  silver 
nitrate,  greater  than  is  sufficient  to  precipitate  the  compound, 
then  titrating  this  with  sulphocyanate  to  learn  what  the  ex- 
cess is.  The  difference  in  cubic  centimeters  between  the 
volumes  of  silver  nitrate  used  and  sulphocyanate  used  shows 
the  amount  of  standard  silver  solution  which  is  taken  up  by 
the  substance  under  investigation.  From  this  its  weight 
is  obtained.  The  end  reaction  is  rather  easier  to  distinguish 
than  when  potassium  chromate  is  the  indicator. 

The  standard  solutions  of  silver  nitrate  and  potassium  sul- 
phocyanate may  be  employed  in  the  estimation  of  any  com- 
pounds of  the  metals  which  can  be  converted  easily  and  without 
loss  into  chlorids.  This  can  be  done  with  the  carbonates, 
hydroxids,  'oxids  and  nitrates  of  potassium,  sodium,  ammo- 


l8o  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

nium,  calcium,  strontium,  barium,  magnesium,  and  some  other 
metals  by  the  action  of  hydrochloric  acid;  also  with  the  chlo- 
rates, which  by  ignition  set  free  oxygen  and  are  changed  to 
chlorids.  Carbonates  may  be  decomposed  with  the  evolu- 
tion of  carbon  dioxid  and  nitrates  with  a  setting  free  of  nitric 
acid.  This  is  accomplished  by  adding  to  a  weighed  or  meas- 
ured amount  of  the  substance  under  investigation  an  excess 
of  the  acid  (after  evaporation  to  dryness  in  case  of  solutions) 
using  concentrated  acid  with  nitrates.  The  excess  of  the 
acid,  that  is,  the  part  which  has  not  united  with  the  metal  to 
form  chlorids,  must  be  completely  driven  off  by  first  evaporat- 
ing to  dryness  on  a  steam-bath,  then  heating  in  an  air-bath 
.at  120°  until  a  piece  of  blue  litmus-paper  laid  across  the  dish 
is  no  longer  reddened.  The  chlorid  is  then  dissolved  in  water 
and  this  solution  titrated  as  before,  using  a  known  fraction 
of  the  liquid  and  making  duplicate  determinations.  After 
the  determination  of  the  chlorin  the  weight  of  the  metal,  and 
hence  of  the  original  compound,  can  be  calculated. 

Practical  Exercises  in  Analysis  by  Precipitation 

1 .  Prepare  decinormal  silver  nitrate  by  one  of  the  methods 
given  above. 

2.  With  this  solution  determine  the  percentage  concentra- 
tion of  a  solution  of  sodium  chlorid.     Use  only  enough  of  the 
potassium  chromate  to  make  the  solution  slightly  yellow,  which 
will  take  but  a  few  drops.     With  a  large  quantity  it  becomes 
more  difficult  to  tell  when  the  silver  chromate  commences  to  be 
permanent.     Make  the  titration  by  yellow  light  or  gas  light  if 
convenient  (not  that  of  an  incandescent  burner,  however). 
One  molecule  of  silver  nitrate  precipitates  oneof  the  chlorid. 

AgN03+NaCl=AgCl+NaN04. 

3.  In  the  same  manner  determine  the  concentration  of  a 
solution  of  potassium  bromid. 


ANALYSIS  BY  PRECIPITATION  l8l 

4.  A  potassium  cyanid  solution  upon  the  addition  of  the 
silver  ion  forms  at  first  no  precipitate,  but  the  soluble  salt 
KAg(CN)2. 

AgN03+KCN  =  KAg(CN)2+HN03. 

When  all  the  potassium  cyanid  has  been  thus  changed  any 
excess  of  silver  nitrate  decomposes  the  double  salt  and  a  pre- 
cipitate appears.  This  reaction  can  be  made  use  of  as  a  quan- 
titative method.  Into  10  c.c.  of  the  potassium  cyanid  solution 
run  from  the  burette  decinormal  silver  nitrate,  stirring 
continually.  The  equation  above  shows  that  as  soon  as  a 
permanent  precipitate  appears,  for  each  molecule  of  the  silver 
nitrate  used  two  of  the  cyanid  have  entered  into  reaction. 
One  c.c.  of  decinormal  silver  nitrate  accordingly  corresponds 
here  to  0.013  grm-  °f  potassium  cyanid,  that  is,  twice  as  much 
as  in  corresponding  reactions  between  other  salts  and  silver 
nitrate,  because  of  the  formation  of  the  double  salt.  In  this 
method  no  indicator  is  necessary.  Ascertain  by  it  the  per- 
centage concentration  of  a  solution  of  the  cyanid. 

5.  With  hydrocyanic  acid  silver  nitrate  gives  the  following 
reaction : 

AgN03+HCN=AgCN+HN03. 

The  nitric  acid  thus  formed  would  interfere  with  the  titra- 
tion  and  must  be  neutralized.  This  can  be  done  by  stirring 
into  the  hydrocyanic  acid  that  is  to  be  titrated,  after  measur- 
ing, enough  of  an  aqueous  suspension  of  magnesia  to  give  an 
alkaline,  milky  mixture.  Then  add  a  few  drops  of  potassium 
chromate  solution  and  the  standard  silver  nitrate  until  the 
red  silver  chromate  does  not  disappear  after  stirring.  Avoid 
inhaling  the  vapor  of  hydrocyanic  acid:  It  should  not  be 
drawn  into  the  pipette  with  the  mouth.  It  is  better  to  weigh 
the  original  portion  than  to  attempt  to  measure  it. 

6.  Prepare  a  decinormal  solution  of  potassium  sulphocya- 
nate  after  the  method  outlined  above  (p.*  178). 


1 82  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

With  the  standard  solutions  of  silver  nitrate  and  sulpho- 
cyanate  ascertauVthe  concentration  of  a  solution  of  ammonium 
chlorid  by  the  method  of  residual  titration.  To  10  c.c.  of  the 
chlorid  solution  add  5  c.c.  of  the  ammonium  ferric  sulphate  in- 
dicator and  5  c.c.  of  dilute  nitric  acid.  From  a  burette  run  in 
a  few  drops  of  decinormal  potassium  sulphocyanate,  then, 
from  another  burette,  enough  decinormal  silver  nitrate  to 
make  the  color  of  the  liquid  a  pure  white.  If  sufficient  has 
been  used,  another  addition  of  the  sulphocyanate  will  give  no 
red  color.  If  it  does,  continue  adding  the  silver  solution  until 
this  condition  is  attained.  The  chlorid  ion  is  then  all  precipi- 
tated as  well  as  the  whole  of  the  sulphocyanate,  and  there  is 
present  an  excess  of  silver  nitrate.  Now  allow  the  sulphocya- 
nate Solution  to  flow  in  slowly  from  the  burette  with  constant 
stirring  until  the  red  color,  which  disappears  at  first,  is  per- 
manent. This  indicates  that  both  the  chlorid  and  sulphocya- 
nate ions  are  exactly  precipitated.  The  difference  between 
the  volumes  of  silver  nitrate  and  .potassium  sulphocyanate 
used  gives  the  amount  of  decinormal  silver  solution  which  was 
required  to  precipitate  the  chlorid  ion.  From  this  the  weight 
of  the  chlorin  can  be  calculated. 

7.  Determine  the  concentration  of  a  calcium  chlorid  solu- 
tion by  making  it  alkaline  with  ammonium  hydroxid,  then 
adding  ammonium  carbonate  solution  as  long  as  precipitation 
occurs  and  heating  to  boiling  for  a  few  minutes.     Filter  and 
wash  with  hot  water  until  the  wash- water  is  no  longer  alkaline. 
Then  place  the  precipitate  with  the  paper  in  a  beaker  with  a 
small  quantity  of  water.     With  methyl-orange  as  an  indicator, 
titrate  the  liquid  with  normal  hydrochloric  acid  until  it 
changes  to  a  permanent  pink  (acid)  color.     Calculate  the 
percentage  of  the  calcium  salt  in  solution. 

8.  Mercuric  chlorid  when  poured  into  a  solution  of  potas- 
sium iodid  forms  a  red  precipitate  (89).     This  dissolves,  as  it 
is  stirred,  as  long  as  no  more  than  one  molecule  of  mercuric 


ANALYSIS  BY  PRECIPITATION  183 

chlorid  is  present  for  four  of  potassium  iodid.  When  there  is 
more  than  this  the  red  precipitate  is  permanent.  The  mer- 
curic solution  must  be  poured  into  the  iodid  and  not  the 
opposite.  By  this  means  the  metal  or  any  mercury  compound 
can  be  quantitatively  determined  provided  it  can  be  first  con- 
verted into  mercuric  chlorid.  A  decinormal  solution  of  potas- 
sium iodid,  made  by  dissolving  the  weighed  crystals  in  water, 
may  be  used  as  the  standard  and  one-fourth  the  molecular 
weight  of  mercuric  chlorid  reckoned  as  present  for  each 
molecule  of  the  iodid  necessary  to  obtain  the  permanent 
precipitate. 

9.  Make  a  determination  of  the  amount  of  ammonium 
nitrate  in  a  solution  of  this  salt  which  also  contains  ammo- 
nium sulphate,  using  the  method  given  above  (page  180)  of 
converting  the  nitrate  into  a  chlorid. 

10.  Neutral  soluble  salts  of  which  the  bases  can  be  com- 
pletely precipitated  by  sodium  carbonate  can  also  be  estimated 
by  the  use  of  standard  solution  of  the  latter,  the  reaction 
being  of  the  following  type: 

Ba+C03"  =  BaC03. 

When  the  reaction  of  the  liquid  changes  to  alkaline  it  indi- 
cates that  the  cation  has  been  precipitated.  Salts  of  barium, 
strontium,  and  calcium  can  be  quantitatively  determined  thus. 
To  ensure  the  certainty  that  they  are  fully  precipitated  the 
liquid  must  be  heated  to  boiling. 

Find  the  concentration  of  a  solution  of  barium  nitrate, 
Ba(NO3)2,  after  the  addition  of  a  few  drops  of  phenolph- 
thalein  solution,  by  titrating  with  normal  sodium  carbonate 
while  the  barium  solution  is  at  the  boiling-point. 

Questions  for  Further  Study  in  Volumetric  Analysis 

What  is  the  relation  of  the  cubic  centimeter  to  the  gramme? 
Of  the  gramme  to  the  liter?  How  many  grammes  in  a  liter 


1 84  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

of  water  ?  Is  this  the  same  for  all  liquids  ?  What  is  the  best 
method  of  preparing  standard  solutions  from  very  hygroscopic 
or  efflorescent  compounds  ?  What  are  some  of  the  most  com- 
monly used  reagents  of  these  classes  ?  If  a  compound  con- 
taining water  of  crystallization  is  to  be  weighed  to  make  up 
such  a  solution  how  can  it  be  determined  whether  the  theoret- 
ical amount  of  water  is  present  or  not  ?  In  acidimetry  which  is 
determined,  the  cation  or  anion?  Which  in  alkalimetry? 
What  would  be  the  best  method  for  preparing  a  standard  solu- 
tion of  ammonium  hydroxid?  Since  both  the  standard  sodi- 
um hydroxid  and  sodium  carbonate  have  an  alkaline  reaction 
and  will  neutralize  acids,  why  is  it  important  to  prevent  the  con- 
version of  the  former  to  the  latter  by  carbon  dioxid  when  it  is 
to  be  used  in  a  standard  solution  ?  Why  cannot  standard  acids 
like  hydrochloric,  sulphuric,  and  nitric  be  prepared  by  accu- 
rately weighing  the  pure  concentrated  acid  and  diluting  this 
with  a  known  weight  of  water  ?  Which  ones  of  the  common 
standard  solutions  are  stable,  and  which  must  be  standardized 
at  each  time  of  using?  What  is  the  nature  of  the  blue  com- 
pound formed  by  starch  and  iodin  andjwhat  would  be  the  effect 
of  heating  the  liquid  which  contains  it?  Are  standard  solu- 
tions of  reducing  agents  permanent  or  not  ?  Why  ?  Are  those 
of  oxidizing  agents  more  or  less  permanent  ?  Under  what  con- 
ditions are  both  best  preserved?  Why  is  ferrous  ammonium 
sulphate  selected  for  the  standard  solution  of  iron  instead  of 
crystallized  ferrous  sulphate  ?  Given  10  grammes  of  dry,  pure, 
sodium  carbonate  and  the  necessary  measuring  flasks,  pipettes, 
etc.,  how  can  decinormal  solutions  of  silver  nitrate  be  pre- 
pared? of  potassium  permanganate?  of  sodium  hydroxid? 
Starting  in  the  same  manner  from  10  grammes  of  pure  crystal- 
lized oxalic  acid,  how  is  it  possible  to  prepare  standard  solutions 
of  hydrochloric  acid?  of  ferrous  sulphate?  of  silver  nitrate? 


PART  III 

APPLIED  ANALYSIS 


CHAPTER  I 

THE  SANITARY  EXAMINATION  OF  WATER 

ALL  natural  waters,  whether  known  as  well,  spring,  river,  or 
cistern  water,  have  been  precipitated  to  the  earth  as  rain  and 
before  they  can  be  used  have  been  brought  into  contact  with 
so  many  soluble  substances  that  they  are  never  chemically 
pure.  The  purest  form  is  rain  water,  but  this  contains  in 
solution  the  gases  of  the  atmosphere  and  often  in  addition 
small  particles  which  floated  in  the  air  in  the  form  of  dust.  If 
it  has  passed  through  or  over  the  soil  it  contains  more  or  less 
of  the  mineral,  vegetable,  or  animal  compounds  with  which  it 
has  been  in  contact.  The  dissolved  gases  are  in  general  harm- 
less and  usually  improve  the  taste  of  the  water  which  without 
them  is  called  flat  or  insipid.  Of  the  other  two  classes  of  im- 
purities— the  mineral  and  organic — the  mineral  are  much  the 
less  objectionable.  Unless  they  are  present  in  large  quantities 
they  produce  little  physiological  effect.  If  the  amount  is 
sufficiently  great  to  affect  the  system,  the  water  is  classed 
with  the  mineral  waters  which  are  so  often  used  on  account  of 
their  therapeutic  value.  Calcium  and  magnesium  salts  (prin- 
cipally the  carbonate  and  sulphate)  give  to  water  the  property 
commonly  called  "  hardness,"  that  is,  it  does  not  dissolve  soap, 
and  consequently  does  not  have  the  soft  feeling  that  is  charac- 
teristic of  soapy  water.  It  is  also  impossible  to  produce  a  per- 
manent foam  or  lather  with  such  hard  water  by  shaking  it  with 

185 


1 86  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

soap  solution  unless  a  large  amount  of  soap  is  added.  Con- 
siderable quantities  of  mineral  matters  are  objectionable 
where  the  water  is  to  be  used  for  making  steam  as  it  sometimes 
deposits  as  " boiler  scale"  and  sometimes  attacks  the  iron  of 
the  boiler.  Compounds  of  lead,  copper,  or  zinc  should  con- 
demn a  water,  since  they  are  poisonous,  but  they  are  rarely 
found  in  it.  The  chlorid  ion  is  a  valuable  indication  of  the 
pollution  of  water.  Sodium  chlorid  is  found  in  soils  where 
this  has  not  occurred,  but  then,  as  a  rule,  only  in  very  small 
quantities.  The  normal  amount  in  any  locality  can  be  deter- 
mined by  testing  the  water  from  a  number  of  sources  in  the 
region,  and  is  practically  a  constant.  As  salt  is  always  present 
abundantly  in  animal  excreta  and  as  it  is  soluble  it  will  be 
taken  up  by  any  water  that  may  come  in  contact  with  it. 

Organic  matter,  vegetable  or  animal,  in  water  may  be 
considered  dangerous,  and  when  abundant  should  condemn 
the  water  for  drinking.  Unfortunately,  the  term  organic 
matter  does  not  mean  one  chemical  compound,  but  rather  a 
complex  mixture  of  varying  nature  so  that  its  identification 
is  often  a  matter  of  some  difficulty.  This  is  especially  true  of 
the  distinction  between  vegetable  and  animal  substances, 
which  are  composed  of  the  same  elements  and  contain  similar 
compounds.  They,  therefore,  respond  to  the  same  test  in 
many  cases.  They  are,  moreover,  when  contained  in  water, 
in  a  state  of  decomposition,  becoming  converted  into  other 
compounds.  Thus  the  nitrogen,  which  is  one  of  the  charac- 
teristic elements  of  this  mixed  organic  matter,  appears  after 
bacterial  disintegration,  first  in  ammonia,  then  in  nitrites,  and 
lastly  in  nitrates.  Since  none  of  the  last  three  are  found  in 
common  soil  or  air,  except  in  traces,  these  may  be  assumed 
to  indicate  that  the  water  contains,  or  has  at  some  time  con- 
tained, vegetable  or  animal  matter. 

Non-living  organic  matter  in  water  may  have  deleterious 
effects  upon  the  animal  system  from  its  own  action  or  that  of 


THE    SANITARY   EXAMINATION    OF    WATER  187 

its  decomposition  products,  or  it  may  serve  as  food  for  micro- 
organisms of  which  the  pathogenic  bacteria  are  the  most  im- 
portant. These  are  thrown  off  in  the  excreta  from  many  in- 
fectious diseases  like  typhoid  and  many  can  live  a  long  time 
in  water  and  communicate  the  disease  by  this  means.  Be- 
cause of  its  frequent  association  with  such  organisms  water 
which  is  contaminated  with  animal  matter  is  regarded  as 
more  dangerous  than  that  which  contains  a  similar  amount  of 
vegetable  products.  Vegetable  and  animal  substances  re- 
main if  water  is  evaporated  and  when  more  highly  heated 
they  are  discolored  and  blackened,  often  giving  rise  to  offen- 
sive or  characteristic  odors.  They  are  oxidized  by  potassium 
permanganate  in  the  presence  of  sulphuric  acid  while  they  are 
in  solution,  the  permanganate  being  at  the  same  time  de- 
colorized. If  the  water  has  been  proved  free  from  nitrites  or 
ferrous  compounds,  the  action  of  which  is  the  same,  this  may 
be  considered  a  test  for  the  presence  of  organic  substances. 

Qualitative  Tests 
Calcium  and  Magnesium 

411.  Make  50  c.c.1  of  water  alkaline  with  ammonium 
hydroxid.  If  there  is  a  precipitate  of  iron  or  aluminum 
hydroxid,  remove  it  by  filtration;  otherwise  to  the  liquid  add 
ammonium  chlorid  and  ammonium  oxalate.  A  fine  white 
precipitate  of  calcium  oxalate  forms  slowly  if  but  little  cal- 
cium is  present.  If  the  precipitate  is  more  than  is  obtained  in 
the  test  for  sulphates  it  indicates  that  a  part  of  the  calcium 
is  probably  in  the  form  of  a  carbonate.  Let  it  settle  in  a 
warm  place,  then  decant  or  filter  and  test  the  filtrate  for  mag- 
nesium by  the  addition  of  sodium  phosphate.  A  white  crystal- 
line precipitate  of  ammonium  magnesium  phosphate  appears. 

1  The  volumes  of  water  directed  to  be  used  in  tests  may  often  be  increased 
if  the  supply  is  large,  thereby  making  the  results  more  accurate.  Except  in 
quantitative  tests  they  need  not  be  accurately  measured. 


1 88  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

• 

Iron 

Examinations  should  be  made  for  both  the  ferrous  and  the 
ferric  forms. 

412.  To  detect  ferrous  compounds  add  a  few  drops  of  a 
dilute   solution   of  potassium  ferricyanid  and  acidify  with 
hydrochloric  acid.     A  blue  color  is  produced.     The  reagent 
does  not  react  in  this  matter  with  the  ferric  ion.     The  ferric 
compounds  may  be  shown  to  be  present,  after  acidifying  with 
hydrochloric  acid,  by  the  production  of  a  red  or  pink  color  by 
the  use  of  potassium  sulphocyanate  which  remains  colorless 
with  the  ferrous  ion. 

Poisonous  Metals 

413.  Lead  is  the  one  which  is  most  frequently  suspected 
from  its  use  in  pipes,  and  occasionally  it  is  desirable  to  test  for 
copper  or  zinc.     With  any  considerable  quantity  of  the  first 
two,  hydrogen  sulphid  gives  a  brown  or  black  color.     It  will 
usually  be  advisable  before  adding  the  sulphid  to  evaporate  to 
about  one-tenth  of  its  bulk,  as  much  as  a  liter  of  water  after 
acidifying  with  hydrochloric  acid,  and  the  use  of  hydrogen 
sulphid  gas  is  here  preferable  to  its  solution,  to  avoid  dilution. 
If  a  black  precipitate  is  obtained,  it  can  be  removed  by  filtra- 
tion and  identified  by  the  method  of  Table  IV.     The  zinc  can 
be  separated  from  the  filtrate  or,  in  absence  of  a  precipitate 
with  hydrogen  sulphid,  from  the  concentrated  water  by  first 
making  the  liquid  slightly  alkaline  with  ammonium  hydroxid, 
then  passing  in  hydrogen  sulphid  and  warming.     It  forms 
white,  voluminous  zinc  sulphid  which  may  be  identified  by 
its  reactions. 

Ammonia 

414.  Ammonia  does  not  occur  in  natural  waters  in  suffi- 
cient abundance  to  render  possible  its  detection  by  its  odor  or 
by  litmus  paper.     Very  minute  quantities,  however,  yield  a 


THE    SANITARY  EXAMINATION   OF   WATER  189 

yellow  to  brown  color  with  Nessler's  reagent.  If  the  amount 
is  considerable  there  is  a  £rown  precipitate,  but  in  ordinary 
waters  the  liquid  remains  clear.  The  result  is  somewhat 
complicated  by  the  presence  of  calcium  and  magnesium  com- 
pounds as  they  are  in  part  precipitated  at  the  same  time. 
The  difficulty  can  be  completely  overcome  by  distilling  a 
part  of  the  water  from  a  retort  and  condensing  the  steam. 
The  pure  ammonia  is  found  in  the  distillate. 

In  the  detection  of  slight  shades  of  color,  or  in  estimating 
their  intensity,  the  desired  result  can  be  best  attained  by 
using  a  test-tube  placed  above  a  piece  of  white  paper  and 
looking  down  into  the  tube. 

Chlorin  (the  Chlorid  Ion) 

415.  Acidify  30  c.c.  of  the  water  with  nitric  acid  and  add  a 
few  drops  of  a  solution  of  silver  nitrate.     An  opalescent  or 
milky  liquid  is  produced  or,  with  large  amounts  of  chlorin, 
a  white  precipitate  of  silver  chlorid.     This  dissolves  in  am- 
monium hydroxid  and  is  reprecipitated  from  this  solution  if 
it  is  acidified  by  nitric  acid.     Ordinarily  there  is  too  little 
chlorin  to  immediately  give  a  curdy  precipitate.     If  this  is 
formed  the  amount  is  excessive. 

Sulphates 

416.  Acidify  30  c.c.  of  water  with  hydrochloric  acid  and 
add  barium  chlorid.     Fine  white  barium  sulphate  precipi- 
tates, settling  slowly,  and  in  very  dilute  solutions  appearing 
only  after  standing  some  minutes.     It  is  insoluble  in  acids. 

Nitrous  Acid  or  Nitrites 

417.  In  the  absence  of  ferric  compounds  the  following 
may  be  used,  but  not  if  ferric  salts  are  present.     To  30  c.c. 
of  water  add  two  or  three  drops  each  of  solutions  of  starch 
and  potassium  iodid,  then  acidify  with  about  ten  drops;;of 


I  go  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

dilute  sulphuric  acid.     Nitrous  acid  sets  iodin  free,   which 
with  the  starch  forms  a  blue  color. 

Ferric  salts  do  not  interfere  with  the  next  two  tests. 

418.  Add  to  30  c.c.  of  water,  1-2  c.c.  of  dilute  sulphuric  acid 
and  i  c.c.  of  a  solution  of  meta-phenylen-diamin  sulphate, 
which  latter  must  be  colorless.     The  appearance  of  a  yellow 
to  brown  color  indicates  the  presence  of  the  nitrite  ion. 

419.  Twenty  c.c.  of  water  with  2  c.c.  of  alpha-amido-naph- 
thalene  (naphthylamin)  acetate  and  2  c.c.  of  sulphanilic  acid 
dissolved  in  acetic  acid  gives  a  pink  or  red  color,  varying  with 
the  amount  of  nitrites  present.     It  should  be  allowed  to  stand 
five  minutes  to  develop  this  fully. 

The  last  two  tests  are  much  more  sensitive  than  the  first, 
so  much  so  that  extreme  care  must  be  used  that  neither  they 
nor  the  vessels  which  contain  the  water  are  contaminated  by 
nitrous  acid  from  the  atmosphere.  The  third  test  will  detect 
one  part  of  nitrous  acid  in  one  hundred  million  parts  of  water. 
On  the  other  hand,  if  the  starch  test  is  distinct,  in  the  ab- 
sence of  iron  compounds,  it  is  probable  that  a  considerable 
amount  of  the  nitrite  ion  is  present. 

Nitric  Acids  or  Nitrates 

If  the  nitrite  ion  has  been  found  by  the  preceding  tests  it 
must  be  removed  before  examining  for  the  nitrate  ion  as  it 
will  give  most  of  the  reactions  for  the  latter.  Evaporation  of 
the  water  to  dryness  after  the  addition  of  a  small  pinch  of  am- 
monium chlorid  decomposes  the  nitrites  but  not  the  nitrates. 

NaNO2+NH4Cl  =  NaCl+  2H2O+N2. 

The  evaporation  can  be  performed  in  a  porcelain  dish.  To 
avoid  the  possibility  of  overheating  at  the  last,  when  it  has 
been  reduced  to  a  small  volume  it  should  be  placed  upon  a 
beaker  of  boiling  water  and  heated  thus  until  it  is  dry.  The 
residue  can  then  be  dissolved  in  pure  water  and  the  solutions 


THE    SANITARY   EXAMINATION    OF    WATER  IQI 

tested  for  nitrates.     They  will  readily  go  into  solution,  al- 
though some  of  the  other  solids  may  remain. 

420.  To  30  c.c.  of  water  add  enough  indigo  solution  to 
impart  a  pale  blue  color,  acidify  with  sulphuric  acid  and  heat 
to  boiling.     If  nitrates  are  present  the  blue  color  is  discharged 
or  diminished. 

421.  In  a  porcelain  dish  containing  2-3  c.c.  of  water  place 
a  small  crystal  of  brucin,  then  add  an  equal  volume  of  con- 
centrated sulphuric  acid,  stirring  gently  with  a  glass  rod. 
A  red  color  indicates  nitric  acid. 

422.  In  the  same  manner  as  the  last  repeat  the  test  using  a 
small  crystal  of  di-phenyl-amin  instead  of  brucin.     Nitric 
acid  or  its  salts  give  a  deep  blue  color.     With  one  part  in 
100,000  it  is  seen  immediately;  when  twice  as  dilute  it  may 
not  appear  for  several  minutes. 

The  test  with  indigo  is  not  nearly  so  sensitive  as  the  last 
two,  but  only  responds  when  a  considerable  amount  of 
nitrates  are  present.  The  concentrated  sulphuric  acid  of  the 
trade,  even  the  so-called  chemically  pure,  frequently  contains 
nitric  acid  and  should  be  proved  to  be  pure  previous  to  using 
in  all  the  above  tests.  If  not  so,  the  nitric  acid  may  be  ex- 
pelled by  boiling  for  half  an  hour,  until  it  remains  colorless 
when  tested  with  diphenyl-amin. 

Hydrogen  Sulphid 

423.  Acidify  slightly  50  c.c.  to  100  c.c.  of  water  with  sul- 
phuric acid  and  boil  for  some  time  in  a  small  flask  in  the  neck 
of  which  is  suspended  a  strip  of  white  filter-paper  previously 
moistened  with  a  solution  of  lead  acetate.     A  yellow  or  brown 
discoloration  of  the  paper  indicates  hydrogen  sulphid. 

Organic  Matter,  Vegetable  and  Animal 

Many  animal  and  vegetable  nitrogenous  compounds  are 
oxidized  by  potassium  permanganate  in  alkaline  solution  with 


1 92  INTRODUCTION   TO    CHEMICAL  ANALYSIS 

a  conversion  of  the  nitrogen  to  ammonia.  This,  from  its 
origin,  is  called  albuminoid  ammonia  and  in  general  indicates 
the  amount  of  such  nitrogenous  compounds  in  the  water. 
Many  of  the  organic  compounds  which  are  found  in  water 
are  readily  oxidized,  even  when  dissolved,  by  such  reagents 
as  potassium  permanganate.  This  loses  its  red  color  at  the 
same  time.  A  few  inorganic  substances  have  also  the  ability 
to  decolorize  permanganate,  those  which  may  be  present  in 
natural  waters  being  sulphids,  nitrites,  and  ferrous  compounds. 
If  these  are  absent  or  in  very  small  amounts,  water  which  de- 
colorizes more  than  a  few  drops  of  a  permanganate  solution 
may  be  considered  to  contain  some  form  of  organic  matter. 

424.  Acidify  50  c.c.  to  100  c.c.  of  the  water  in  a  beaker  or 
porcelain  dish  with  5-10  c.c.  of  dilute  sulphuric  acid,  heat  to 
boiling,  and  while  hot  drop  in  the  permanganate  solution 
until  a  permanent  pink  tint  remains.     Compare  the  result  with 
that  obtained  in  the  same  manner  from  a  similar  quantity  of 
distilled  water. 

In  addition  to  the  decomposition  products  already  dis- 
cussed there  are  produced  in  water  by  bacterial  processes 
small  amounts  of  a  number  of  compounds  of  which  the  most 
common  are  indol,  phenol,  and  their  derivatives.  Griess  has 
proposed  a  method  of  detecting  these  by  the  intensely  yellow 
color  which  they  yield  with  diazo  compounds.  They  are  very 
abundant  in  water  which  is  polluted  with  decaying  animal 
matter  or  animal  excreta.  They  may,  however,  be  formed  by 
vegetable  matter,  their  source  being  the  protein  compounds 
which  are  found  in  both,  but  less  abundantly  in  the  latter. 

425.  Make  alkaline  with  sodium  hydroxid  25  c.c.  of  water 
and  add  a  few  drops  of  a  freshly  prepared  dilute  solution 
of  para-di-azo-benzene  sulphonic  acid.     The  liquid  becomes 
yellow  if  the  water  contained  decomposing  animal  matter 
or  sometimes  with  large  amounts  of  vegetable  matter.     The 
reagent  cannot  be  preserved  in  solution,  but  must  be  dissolved 


THE    SANITARY   EXAMINATION    OF    WATER  1 93 

only  when  it  is  to  be  used.  With  the  alkali  and  distilled  water 
there  should  be  thus  produced  no  color,  although  this  may 
occur  if  the  sodium  hydroxid  is  impure.  A  test  should  there- 
fore first  be  made  with  distilled  water  to  ascertain  the  purity 
of  the  reagents.  This  is  Griess'  test. 

Quantitative  Determinations 
Total  Solids 

The  amount  of  solid  matter  dissolved  in  a  water  is  some- 
times of  importance.  This  can  be  ascertained  by  evaporat- 
ing to  dryness  on  a  steam-bath  a  definite  volume  in  a  weighed 
dish  and,  after  cooling,  weighing  the  dish  with  the  residue. 
With  a  sensitive  balance  100  c.c.  of  water  is  enough  in  most 
cases,  but  if  the  balance  will  not  weigh  to  a  milligramme,  more 
may  be  taken.  A  large  amount  of  dissolved  solids  is  indica- 
tive of  contamination,  although  not  conclusive.  Mineral 
waters  are,  of  course,  an  exception.  The  nature  of  the  resi- 
due is  indicated  if  it  is  more  strongly  heated  after  weighing. 
This  is  best  done  in  a  platinum  dish,  but  may  be  accomplished 
in  one  of  thin  porcelain.  Much  organic  matter  will  turn 
brown  or  black  and  may  give  a  more  or  less  characteristic  odor. 
Mineral  substances  are  not  discolored,  with  the  exception  of 
iron  compounds  which  may  be  brown.  After  the  organic 
matter  has  burned  off  they  can  be  dissolved  in  hydrochloric 
acid  (except  silica)  and  identified  by  their  qualitative  reactions. 

Ammonia 

The  exact  determination  of  ammonia  is  effected  by  distil- 
ling it  from  a  large  volume  of  water  in  a  retort.  The  con- 
densed steam  contains  the  ammonia.  By  comparing  the 
color  produced  in  it  by  Nessler's  reagent  with  that  obtained 
by  the  same  reagent  in  pure  water  to  which  has  been  added 
a  known  weight  of  ammonia,  its  amount  can  be  estimated. 

13 


1 94  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

There  are  required  a  standard  solution  of  ammonia  and  an 
alkaline  permanganate  solution.  The  standard  ammonia 
solution  is  prepared  by  dissolving  3.82  grms.  of  pure  am- 
monium chlorid  in  water  and  diluting  to  a  liter.  This  con- 
tains a  milligramme  of  ammonia  (NH3)  in  a  cubic  centimeter. 
For  actual  use  it  is  better  to  dilute  this  100  times.  In  the 
preparation  of  these  solutions  only  ammonia-free  water 
should  be  used — that  is,  water  from  which  the  ammonia  has 
been  removed  by  distillation  or  boiling. 

JThe  alkaline  permanganate  solution  is  made  by  dissolving 
8  grms.  of  potassium  permanganate  and  200  grams  of  po- 
tassium hydroxid  in  1,200  to  1,500  c.c.  of  water  and  distilling 
until  the  remaining  liquid  is  a  liter  and  the  distillate  shows 
no  test  for  ammonia. 

426.  Connect  with  a  condenser  a  700-800  c.c.  glass-stop- 
pered retort  which  has  the  neck  so  bent  that  by  inclining  it 
upward  drops  spattered  from  the  boiling  liquid  will  run 
back  into  the  retort  and  only  the  vapor  will  pass  into  the  con- 
denser. (The  retort  and  condenser  are  connected  as  in  Fig. 
19.)  Use  500  c.c.  of  water  made  alkaline  by  3-4  c.c.  of  sodium 
carbonate  solution  freed  from  ammonia  by  previous  boiling. 
A  few  pieces  of  pumice  stone  ignited  and  dropped  while  hot 
into  the  retort  makes  the  distillation  proceed  more  quietly. 

Collect  the  distillate  in  large  test-tubes  or  flat-bottomed 
Nessler  tubes  in  portions  of  50  c.c.  each,  adding  to  each  2  c.c. 
of  Nessler's  reagent.  Drop  from  a  burette  into  ammonia- 
free  water  enough  of  the  standard  ammonia  solution  to  give 
the  same  shade  with  Nessler's  reagent  when,  after  standing 
five  minutes,  it  has  been  diluted  to  50  c.c.  In  this  manner 
the  amount  of  ammonia  in  each  portion  is  learned.  The 
free  ammonia  should  be  all  expelled  with  the  first  200  c.c. 
Then  discontinue  the  heating  and  pour  in  50  c.c.  of  the  alka- 
line permanganate  solution;  continue  the  distillation  and 
Nesslerizing  as  before  as  long  as  ammonia  is  driven  off.  This 


THE    SANITARY   EXAMINATION    OF    WATER  IQ5 

portion  is  the  albuminoid  ammonia  and  represents  the  un- 
decomposed  nitrogenous  organic  matter. 

Chlorin  (the  Chlorid  Ion) 

427.  To  determine  the  quantity  of  chlorin  use  a  standard 
solution  of  silver  nitrate  with  normal  potassium  chromate  for 
the  indicator,  as  described  on  page  180.  For  convenience  in 
calculating  results  instead  of  the  usual  concentration  the 
silver  solution  may  be  made  to  contain  4.79  grammes  of  silver 
nitrate  per  liter.  One  cubic  centimeter  of  this  will  precipitate 
one  milligramme  of  chlorin.  Of  ordinary  well-water  100  c.c. 
may  be  used.  In  this  case  each  cubic  centimeter  of  the  silver 
solution  necessary  indicates  one  part  of  chlorin  in  100,000  of 
water.  With  water  of  great  purity  it  may  be  necessary  to 
concentrate  a  larger  volume,  such  as  a  liter,  before  titrating. 

Nitrites  (the  Nitrite  Ion) 

When  solutions  of  nitrites  are  treated  with  sulphanilic  acid 
and  naphthylamin  an  intense  rose-red  color  appears  varying 
with  the  amount  of  the  nitrite  ion  present.  By  comparing  it 
with  a  solution  of  nitrite  of  known  strength  it  is  not  difficult 
to  judge  the  amount  of  the  latter. 

The  reagents  used  are: 

Sulphanilic  acid,  0.5  grm.  dissolved  in  150  c.c.  of  3  per  cent, 
acetic  acid,  and 

Alpha-nap  hthylamin,  o.i  grm.  of  the  solid  boiled  with  20 
c.c.  of  water  and  filtered  hot  through  a  small  filter,  previously 
well  washed  to  remove  nitrites,  then  the  filtrate  diluted  with 
180  c.c.  of  3  per  cent,  acetic  acid. 

Standard  sodium  nitrite,  0.049  grm-  °f  pure  sodium  nitrite 
dissolved  in  one  liter  of  water.1  One  cubic  centimeter  of  this 
contains  o.oi  milligramme  of  nitrogen.  If  it  cannot  be  ob- 

1  Since  this  would  be  difficult  to  weigh  accurately  it  is  better  to  dissolve 
0.49  grm.  in  100  c.c.  and  dilute  10  c.c.  of  this  to  1000  c.c.  before  using. 


196  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

tained  pure  in  the  dry  state  it  can  be  made  by  dissolving  the 
corresponding  weight  (0.113  grm-)  of  silver  nitrite  in  a  small 
volume  of  water  and  adding  a  solution  of  pure  sodium  chlorid 
as  long  as  a  precipitate  forms,  then  diluting  to  the  desired 
amount.  It  must  be  kept  in  the  dark.  The  reagent  bottles 
and  all  vessels  used  in  the  test  must  be  rinsed  with  nitrite- 
free  water. 

428.  To  exactly  10  c.c.  of  the  water  in  a  test-tube  add  i  c.c. 
each  of  the  solutions  of  sulphanilic  acid  and  naphthylamin 
and  allow  to  stand  five  minutes.     If  it  is  pink,  in  another 
test-tube  dilute  i  c.c.  of  the  standard  nitrite  solution  to  10  c.c. 
and  test  in  the  same  manner.     Dilute  the  darker  of  the  two 
until  equal  volumes  are  the 'same  shade  and  from  the  amount 
of  dilution  calculate  the  amount  of  nitrite  in  the  water  as  com- 
pared with  that  of  the  standard,  and  from  this,  the  absolute 
weight  of  N  in  nitrites  in  10  c.c.  and  in  a  liter.     Reduce  this 
to  parts  per  million  remembering  that  one  milligramme  per 
liter  equals  one  part  per  million. 

Nitrates  (the  Nitrate  Ion) 

The  nitrate  ion  can  be  quantitatively  determined  by  its 
action  upon  phenolsulphonic  acid,  which  is  changed  into  picric 
acid,  the  ammonium  salt  of  which  has  an  intensely  yellow 
color.  Besides  a  standard  solution  of  a  nitrate  for  compari- 
son but  one  other  is  needed. 

Phenolsulphonic  acid  is  prepared  by  heating  for  an  hour  or 
two  on  the  water-bath  25  grams  of  phenol  (carbolic  acid)  with 
a  mixture  of  150  c.c.  of  concentrated  sulphuric  acid  and  75 
c.c.  of  fuming  sulphuric  acid,  stirring  occasionally.  The 
solution  keeps  well.  If  it  in  time  crystallizes  it  can  be  redis- 
solved  by  warming  gently. 

Standard  potassium  nitrate  can  be  made  by  dissolving  of 
this  salt  0.722  grm.  in  a  liter  of  water.  One  c.c.  of  this 
contains  o.oooi  grm.  of  nitrogen. 

429.  Evaporate  to  dryness  on  a  steam-bath  a  definite  vol- 


THE    SANITARY   EXAMINATION    OF    WATER  197 

ume  of  the  water — 25  to  100  c.c.  according  to  whether  the 
qualitative  test  showed  much  or  little  nitrate  present.  If 
much  chlorin  has  been  found  it  is  best  to  remove  it  by  precipita- 
tion with  nitrate-free  silver  sulphate  and  filtration.  To  the 
residue  add  i  c.c.  of  phenolsulphonic  acid  and,  after  thoroughly 
mixing,  let  it  stand  two  or  three  minutes,  then  add  20  c.c.  of 
water  and  make  alkaline  with  ammonia.  Treat  i  c.c.  of  the 
standard  solution  in  the  same  manner  and  compare  the  colors 
in  large  test-tubes,  diluting  until  they  are  the  same  and  com- 
puting the  amount  of  nitrogen  in  the  nitrates  in  parts  per  mil- 
lion as  with  nitrites  above.  In  the  whole  process  the  distilled 
water  and  sulphuric  acid  used  must  be  known  to  be  free  from 
nitric  acid  and  its  salts,  as  discussed  under  the  qualitative  tests. 

Organic  Matter 

The  amount  of  this  can  be  only  indirectly  determined  be- 
cause of  the  variable  nature  of  that  which  may  be  present. 
The  degree  of  discoloration  or  charring,  also  the  odor,  of  the 
heated  residue  after  evaporation  may  give  some  indication  of 
it.  Other  methods  are  to  determine  the  exact  quantity  of 
organic  carbon  or  nitrogen  (as  in  426)  in  the  water  and  to  use 
them  as  measures  of  the  vegetable  and  animal  matter.  In- 
stead of  these  it  is  more  convenient  to  find  the  amount  of 
oxidizable  matter  and,  deducting  the  nitrites  and  ferrous 
compounds,  also  the  sulphids,  if  they  are  present,  to  regard 
the  remainder  as  organic.  A  suitable  oxidizing  agent  for 
this  purpose  is  a  solution  of  potassium  permanganate  which 
readily  gives  up  its  oxygen  in  an  acid  liquid,  becoming  at  the 
same  time  colorless. 

430.  A  rough  estimate  can  be  made  of  the  amount  of  oxi- 
dizable matter  by  acidifying  50  or  100  c.c.  of  the  water  with  10 
c.c.  of  dilute  sulphuric  acid,  heating  to  boiling,  then  counting 
the  number  of  drops  of  centinormal  potassium  permanganate 
necessary  to  color  the  boiling  liquid  permanently  pink.  The 
reagent  can  be  dropped  slowly  if  the  stopper  is  held  loosely 


198  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

without  removing  it  from  the  bottle.  For  comparison  the 
same  test  may  be  tried  on  an  equal  volume  of  distilled  or 
other  pure  water. 

431.  For  more  accurate  results  a  centinormal  solution 
(0.316  grm.  per  liter)  of  potassium  permanganate  may  be  used 
with  a  corresponding  one  of  oxalic  acid,  so  that  one  cubic 
centimeter  of  the  latter  shall  decolorize  exactly  one  of  the 
former.  To  100  c.c.  of  the  water  in  a  250  c.c.  flask  add  20  c.c. 
of  dilute  sulphuric  acid,  then  from  a  burette,  10  c.c.  of  the  per- 
manganate or  enough  to  give  permanent  red  color  to  the  water 
after  it  has  been  boiled  10  minutes.  This  should  be  accu- 
rately measured.  When  it  has  been  boiled  10  minutes  add  from 
another  burette  a  volume  of  centinormal  oxalic  acid  equal  to 
that  of  the  permanganate  used,  when  the  red  will  disappear. 
Now  drop  in  more  permanganate  from  the  first  burette,  while 
stirring,  until  a  permanent  pink  color  is  produced.  The 
volume  of  permanganate  last  added  corresponds  to  that  de- 
stroyed by  the  oxidizable  substance  in  the  water.  Ten  cubic 
centimeters  contain  0.08  milligramme  of  available  oxygen, 
and  the  final  results  should  be  expressed  in  terms  of  this. 

Owing  to  the  fact  that  there  is  a  great  variation  in  the  ease 
of  oxidation  of  different  kinds  of  organic  matter  as  well  as  in 
their  content  of  carbon  and  nitrogen,  the  results  obtained 
should  be  regarded  as  only  approximate  although  they  are 
often  of  great  value  in  arriving  at  a  conclusion  as  to  the 
purity  of  the  water.  For  the  properties  of  the  standard  solu- 
tions of  permanganate  and  oxalic  acid  reference  should  be 
made  to  their  description  under  Oxidimetry. 

Interpretation  of  Results 

Calcium,  magnesium  and  iron  are  comparatively  unim- 
portant in  water  from  the  sanitary  standpoint,  although  when 
excessive  they  greatly  lessen  its  value  for  domestic  purposes, 


THE    SANITARY   EXAMINATION   OF    WATER  199 

the  first  two  because  they  render  it  hard,  and  the  last  because 
it  deposits  as  a  brown  ferric  hydroxid.  They  all  more  or  less 
modify  the  taste.  The  sulphate  ion,  as  found  in  sulphates,  if 
united  with  the  calcium  makes  the  water  permanently 
"hard,"  that  is,  this  compound  is  not  precipitated  by  boiling, 
as  is  the  carbonate.  Small  amounts  of  sulphates  are  other- 
wise not  of  much  consequence,  although  considerable  quanti- 
ties of  sulphates  cause  the  water  to  act  as  a  laxative. 

It  is  only  exceptionally  that  ammonia  is  absent  from  natural 
waters.  It  signifies  that,  probably  by  bacterial  action  nitroge- 
nous organic  compounds,  vegetable  or  animal,  have  been 
decomposed,  the  latter  yielding  more  than  the  former.  When 
nitrites  are  found,  bacterial  action  upon  these  compounds  is  in 
progress,  although  the  bacteria  are  not  necessarily  of  the 
pathogenic  varieties.  Nitrates  are  the  final  products  of  such 
action,  unless  they  are  again  reduced  to  nitrites.  When 
found  alone  they  indicate  some  past  pollution  by  nitrogenous 
substance — animal  or  vegetable.  Chlorin  occurs  rather 
abundantly  in  the  soil  near  the  sea  and  in  parts  of  the  country 
where  there  are  salt  deposits  or  brines.  If  then  loses  much  of 
its  value  as  an  indication  of  contamination.  Since,  however, 
water  drawn  from  the  same  stratum  in  the  same  section  has 
practically  the  same  composition  unless  it  is  contaminated,  a 
comparison  of  the  amount  of  chlorin  in  the  water  of  one  well 
with  that  of  surrounding  ones  will  often  show  whether  it  is 
excessive  or  not.  Neglecting  that  of  the  soil,  it  is  an  evidence 
that  pollution  has  occurred  from  animal  matter,  probably 
sewage.  This  is  of  course  pollution  of  the  most  dangerous 
kind;  still  if  chlorin  alone  is  found  it  may  be  that  the  accom- 
panying organic  matter  and  bacteria  have  been  filtered  out  by 
the  soil,  or  oxidized  by  the  air,  and  that  the  water  will  not  com- 
municate disease;  in  other  words  it  is  an  evidence  of  past  con- 
tamination. The  chlorid  ion  in  itself  has  no  poisonous  action, 
but  its  presence  is  suspicious.  Hydrogen  sulphid  can  be  pro- 


200  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

duced  by  the  breaking  down  of  organic  sulphur  compounds 
like  the  proteins,  or  by  the  action  of  decomposing  organic 
matter  on  the  sulphates,  whereby  the  latter  are  reduced. 
Hence  it  usually  shows  that  considerable  organic  matter  is, 
or  has  been,  contained  in  the  water. 

Decolorization  of  the  potassium  permanganate  solution,  in 
the  absence  of  reducing  agents  like  nitrites,  sulphids,  or  ferrous 
compounds,  indicates  organic  matter,  either  animal  or  vege- 
table. Para-diazo-benzene-sulphonic  acid  in  alkaline  solu- 
tion (Griess'  test)  shows  the  presence  of  compounds  such  as 
result  from  the  putrefaction  of  albuminous  matters.  These 
are  abundant  in  the  excreta  and  decaying  animal  tissues,  but 
may  be  produced  also  from  similar  changes  in  the  vegetable 
albuminous  substances.  In  the  last  case  they  are  only  found 
in  comparatively  small  amounts. 

We  might  then  expect  to  find  in  water  containing  fresh 
vegetable  matter  a  little  ammonia  and  some  reduction  of  the 
permanganate  solution,  but  no  reaction  to  Griess'  test  and  an 
absence  of  more  than  traces  of  chlorin,  nitrites,  and  nitrates. 
As  the  organic  matter  decays  the  ammonia  increases  and  the 
reduction  of  the  permanganate  diminishes  until  there  may  be 
none.  There  will  finally  be  a  conversion  of  most  of  the 
nitrogen  into  nitrates  with  possibly  a  slight  amount  of  nitrites. 
It  is  seldom  that  all  the  ammonia  disappears. 

Animal  matter  (excreta  or  tissues),  while  decomposition  is  in 
progress,  reduces  the  permanganate  solution  and  gives  Griess' 
reaction  if  in  sufficient  quantity.  Chlorin  and  ammonia  are 
then  abundant,  nitrites  are  often  present  and  sometimes 
nitrates.  As  the  decomposition  progresses  the  permanganate 
and  Griess'  test  give  less  marked  results  and  the  nitrates  in- 
crease. The  quantity  of  chlorin  is  not  affected. 

The  source  of  the  water  will  also  somewhat  modify  the  con- 
clusions to  be  drawn  from  the  analysis.  Thus  deep  well  water 
often  contains  much  chlorin  from  the  rock  formations  with 


THE    SANITARY   EXAMINATION    OF    WATER  2OI 

which  it  comes  into  contact.  More  ammonia  may  be  allow- 
able here  than  ?n  surface  or  shallow  well  waters  since  it  has  no 
means  of  escape  and  the  conditions  are  unfavorable  for  bac- 
terial life.  It  is  produced  here  by  the  reduction  of  nitrates 
through  the  action  of  organic  matter.  Nitrates  are  occasion- 
ally found  plentifully  in  deep  waters,  being  derived  from  the 
remains  of  fossil  organisms.  This  is  especially  seen  in  cretaceous 
rocks.  On  the  other  hand,  organic  matter  is  usually  absent 
or  in  traces.  Rain  water  absorbs  much  ammonia  and  some- 
times nitrites  from  the  air  and  these  then  lose  their  customary 
significance,  being  no  longer  indicative  of  organic  decompo- 
sition. In  the  water  from  deep  wells  (artesian)  also  nitrites 
may  be  formed  through  the  reduction  of  nitrates  by  organic 
matter  without  the  aid  of  bacteria  and  then  are  unimportant. 

Standards  of  Comparison 

From  what  has  been  said  regarding  the  modifications  which 
surrounding  conditions  produce  in  the  composition  of  water 
it  is  obvious  that  we  cannot  establish  any  absolute  standards 
which  shall  apply  to  that  from  all  sources.  Nevertheless,  it  is 
a  matter  of  great  convenience  to  be  able  to  compare  the  com- 
position of  one  under  observation  with  that  of  others,  the 
nature  of  which  has  been  determined.  Such  figures  must  be 
used  carefully  with  a  full  understanding  of  the  influences  that 
may  change  them.  Knowing,  then,  that  the  standards  rep- 
resent only  the  amounts  of  impurities  which  have  been  found 
to  be  present  in  other  cases,  we  can  employ  them  as  aids  to 
the  formation  of  an  opinion  as  to  the  significance  of  our 
analyses. 

The  following  may  represent  the  composition  of  some  varie- 
ties of  drinking-water  in  which  the  figures  represent  parts  per 
million. 


2O2 


INTRODUCTION   TO   CHEMICAL   ANALYSIS 


Good, 
less 
than 

Suspi- 
cious. 

Bad, 
over 

Milligrams 
per  Liter. 

=  Parts  per  Million. 

Total  solids, 

Soo 

500-700 

800 

Chlorin, 
Nitrogen  in  nitrites, 
Nitrogen  in  nitrates, 
Free  ammonia, 

5 

0 

0.5 
0.03 

10-15 

0.02 
I  .  O—  2  .  0 
O.O5-O.I 

25 
0.04 
5 
0.15 

NaCl,  2.3 
NaN02,o.i 
KNOs,  14.0 
NEUC1,  0.3 

Chlorin,  15. 
Nitrogen  in  nitrites,  0.02. 
Nitrogen  in  nitrates,  2.0. 
Ammonia,  o.i. 

Albuminoid    ammo- 

nia, 

O.I 

O.I-O.I5 

0.2 

Cubic  centimeters  of 

—  KMn04  reduc'd 

100 

1-5 

3-4 

8 

by  100  c.c.  of  water, 

Solutions  made  up  of  the  above  concentrations  by  the  aid 
of  the  last  two  columns,  may  serve  for  comparison  so  that  the 
more  accurate  quantitative  methods  of  determination  can 
at  times,  be  dispensed  with. 

Practical  Exercises  in  Water  Analysis 

Make  chemical  examinations  of  samples  of  water  furnished 
by  the  instructors.  Hand  in  written  reports  giving  the  pres- 
ence or  absence,  also  comparative  amounts  (small,  moderate, 
or  large)  of  the  following: 

1.  Calcium. 

2.  Magnesium. 

3.  Iron,  ferrous  and  ferric. 

4.  Lead,  copper  or  zinc. 

5.  Ammonia. 

6.  Chlorin. 

7.  Sulphates. 

8.  Nitrites. 

9.  Nitrates. 

10.  Organic  matter  by  general  tests. 

Give  an  opinion  as  to  the  past  history  of  the  water.  If  it 
contains  organic  matter,  is  this  vegetable,  decaying  animal 
or  sewage?  Is  the  water  suitable  for  drinking,  to  be  con- 
demned, or  merely  suspicious  ?  Is  it  suitable  for  other  pur- 
poses than  for  drinking?  Give  reasons  for  your  conclusions. 


THE    DETECTION    OF    POISONS  203 


CHAPTER  II 

THE  DETECTION  OF  POISONS 

THE  testing  of  substances  for  the  detection  and  identifica- 
tion of  poisons,  although  it  may  be  properly  considered  in  a 
separate  chapter  from  general  analytical  chemistry,  is  not  a 
different  branch  of  chemistry,  but  makes  use  of  many  of  the 
principles  and  operations  already  considered. 

Its  importance  lies  in  the  value  of  the  evidence  of  the 
chemist  as  a  proof  of  attempted  or  accomplished  poisoning, 
and  in  the  fact  that  this  is  often  the  strongest,  and  sometimes 
the  only,  evidence  obtainable.  It  differs  from  the  more  com- 
mon methods  of  analysis  in  the  nature  of  the  materials  which 
are  investigated,  these  being  often  complex  mixtures  of  organic 
and  inorganic  compounds,  such  as  medicines,  foods,  vomited 
matter,  saliva,  urine,  blood,  or  animal  tissues.  In  these  the 
poison  sought  is  ordinarily  comparatively  very  small  in 
amount.  Not  only  must  the  presence  of  a  poisonous  element 
be  shown,  but  that  it  is  in  a  combination  dangerous  to  the  ani- 
mal body.  For  example,  most  sulphates  are  harmless,  whereas 
the  acid  from  which  they  are  derived  is  a  violent  poison.  It  is, 
moreover,  frequently  desirable  to  determine  not  only  the  kind 
but  the  amount  of  such  compounds  in  the  substance  analyzed. 

Some  of  the  difficulties  encountered  in  testing  for  poisons 
are  that  their  minute  amount  renders  many  of  the  common 
reagents  ineffective,  that  the  presence  of  other  compounds, 
especially  organic  matter,  interferes  with  their  giving  the  ex- 
pected reaction,  that  their  separation  from  such  compounds 
and  purification  requires  much  time  and  care,  that  during 
the  necessary  operations  they  may  volatilize,  as  does  prussic 


204  INTRODUCTION    TO    CHEMICAL   ANALYSIS 

acid,  or  undergo  decomposition,  like  the  alkaloids,  or  become 
converted  into  a  harmless  form  as  may  be  phosphorus.  Still, 
for  many  of  them,  the  means  of  identification,  if  properly 
carried  out,  are  as  effective  as  any  analytical  processes. 

Since  but  a  limited  quantity  of  the  substance  is  at  the  dis- 
posal of  the  analyst  the  tests  should  be  applied  at  first  to  only 
a  portion  in  order  to  have  a  reserve  in  case  of  accidents,  or  for 
further  investigation,  and  it  is  better  to  note  the  weight  of  the 
whole  and  of  the  part  used  in  order  to  be  able  to  calculate  the 
total  amount  present  from  the  fraction  that  may  be  found  in 
that  part.  To  avoid  a  waste  of  material  the  most  character- 
istic and  reliable  tests  should  be  tried  first.  If  possible  the 
isolated  poison  or  compounds  of  it  should  be  preserved  in  such 
a  way  as  to  be  submitted  in  a  possible  future  trial. 

The  question  of  the  purity  of  the  reagents  is  of  the  greatest 
importance.  The  different  grades  of  "purified,"  "pure," 
"chemically  pure,"  and  "absolutely  chemically  pure"  do  not 
necessarily  guarantee  the  same  quality  when  they  are  made  by 
different  manufacturers.  It  will  be  necessary  to  prove  their 
purity  or  to  make  with  them  a  blank  test — that  is  to  go  through 
the  operation  with  the  chemicals  alone  without  the  addition  of 
the  suspected  substance,  when,  if  they  are  of  satisfactory  qual- 
ity, the  result  should  be  a  negative  one.  The  reagents  in 
common  use  may  suffice  for  the  usual  methods  of  testing  and 
fail  to  give  satisfaction  when  a  hundred  or  a  thousand  times 
the  ordinary  quantity  of  them  must  be  employed.  All  other 
materials,  such  as  filter-paper,  rubber  tubing,  or  apparatus, 
must  also  be  known  to  be  free  from  objectionable  substances. 

There  follows  a  list  of  the  reagents  of  most  frequent  use  in  tox- 
icological  investigations,  with  a  discussion  of  their  impurities. 

Water 

This  should  of  course  be  distilled,  but  as  occasionally 
such  water  contains  traces  of  the  heavy  metals  it  should  be 


THE    DETECTION    OF    POISONS  205 

tested  by  evaporating  a  liter  to  dry  ness.  There  must  be  left 
only  a  minute  residue  and  this,  after  dissolving  in  5-10  c.c. 
of  water  must  give  no  precipitate  with  hydrogen  sulphid  or 
ammonium  sulphid.  Distillation  is  best  conducted  in  a  tin- 
lined  vessel  and  the  steam  condensed  in  a  tin  pipe. 

Hydrochloric  Acid 

This  should,  after  dilution  give  no  precipitate  upon  satura- 
tion with  hydrogen  sulphid  or,  after  neutralization,  with  am- 
monium sulphid .  Its  most  dangerous  contamination  is  arsenic 
which  is  not  infrequent.  Very  large  amounts  of  the  acid  are 
used  in  the  destruction  of  organic  matter  and  hence  a  com- 
paratively large  quantity  of  arsenic  may  be  introduced  even 
if  but  a  small  fraction  of  a  per  cent,  is  contained  in  the  acid. 
It  cannot  be  regarded  as  of  satisfactory  purity  unless  after 
about  a  liter  is  concentrated  to  a  small  volume  (10  to  20  c.c.) 
by  evaporation  in  a  well-ventilated  hood  and  5  c.c.  of  the  re- 
maining liquid  is  tested  with  Gutzeit's  test  (109)  no  yellow 
color  with  brownish-black  margin  appears  on  the  paper  within 
half  an  hour.  It  may  be  further  tested  by  Bettendorf's 
(in),  Reinsch's  (108),  or  Marsh's  (112)  tests. 

Sulphuric  Acid 

This  is  not  used  in  such  large  amounts  as  the  hydrochloric, 
and  there  is  consequently  less  danger  of  introducing  impuri- 
ties with  it.  Its  principal  use  is  in  Marsh's  test  for  arsenic. 
It  occasionally  contains  lead  as  well  as  arsenic.  The  lead 
will  be  shown  by  the  yellow  or  brown  color  produced  by 
saturating  the  dilute  solution  with  hydrogen  sulphid  gas. 
For  arsenic  it  can  be  tested,  after  diluting  with  four  times  its 
volume  of  water,  by  Gutzeit's  (109),  Reinsch's  (108),  or 
Marsh's  (112)  tests.  With  the  latter  200  c.c.  of  the  acid,  gen- 
erating about  a  liter  of  hydrogen  in  fifteen  minutes  should 
show  no  mirror  in  the  tube  after  half  an  hour's  heating. 


206  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

When  used  in  the  alkaloidal  tests  it  must  be  absolutely  free 
from  nitric  or  nitrous  acid.  Hence  it  should  give  no  color 
with  brucin  (421)  or  di-phenyl-amin  (422).  These  com- 
pounds if  present  may  be  removed  by  boiling  as  in  422. 

Hydrogen  Sulphid 

The  materials  commonly  used  in  the  preparation  of  hydro- 
gen sulphid  frequently  contain  arsenic.  Either  those  which 
are  free  from  arsenic  should  be  employed  or  the  gas  must  be 
purified.  Passing  it  through  a  long  tube  over  crystals  of 
iodin  forms  arsenic  iodid  which  remains  in  the  tube. 

Ammonium  Sulphid 

This  must  be  prepared  from  arsenic-free  hydrogen  sul- 
phid, made  in  the  above  manner. 

Ammonium  Hydroxid  and  Sodium  Hydroxid 

Arsenic  and  the  heavy  metals  should  be  absent.  When 
acidified  with  pure  sulphuric  acid  no  results  should  be  ob- 
tained from  Marsh's  test,  from  the  addition  of  hydrogen  sul- 
phid, nor  after  making  alkaline  by  ammonium  sulphid. 

Potassium  Chlorate 

This  should  be  free  from  arsenic  and  the  heavy  metals. 
Dissolve  50  grammes  in  water  and  add  pure  hydrochloric 
acid  as  long  as  chlorin  is  evolved.  Test  half  the  solution  by 
Marsh's  test  (112)  and  the  rest  by  hydrogen  sulphid  (Table 
IV)  and  ammonium  sulphid  (Table  II). 

Ethyl  Alcohol 

Metals  and  other  basic  substances  must  be  absent.  By 
distilling  over  tartaric  acid  they  can  be  separated  if  small 
amounts  are  present. 


THE    DETECTION    OF   POISONS  207 

Amyl  Alcohol 

It  should  leave  no  residue  when  evaporated,  should  boil  at 
I3i°-i32°;  and  when  shaken  with  dilute  sulphuric  acid  should 
give  to  the  latter  nothing  which  gives  the  general  alkaloidal 
reactions  (page  139). 

Benzene 

This  should  have  the  boiling-point  of  79°  and  otherwise 
answer  the  requirements  of  amyl  alcohol. 

Chloroform 
Can  be  tested  as  the  last.     It  should  boil  at  61°. 

Petroleum  Ether 

The  boiling-point  should  not  be  above  60°.  No  residue 
should  remain  when  it  is  allowed  to  evaporate  on  a  watch 
glass. 

The  filter-paper  employed  should  be  white  and  the  ash  re- 
maining after  it  is  burned  should  be  little  and  free  from  bar- 
ium and  the  heavy  metals.  When  the  solutions  must  come  in 
contact  with  the  rubber  tubing  this  should  contain  none  of 
the  poisonous  metals.  In  the  white  tubing  lead  and  zinc  are 
often  found,  and  in  the  red,  antimony.  Black  rubber  is  the 
preferable  kind.  Porcelain  vessels  are  less  liable  to  break 
than  glass  ones,  also  to  attack  by  reagents,  but  they  should 
not  be  used  after  the  enamel  is  in  any  degree  injured,  and  the 
enamel  must  contain  no  lead. 

Outline  of  the  Plan  of  Analysis  for  Poisons 

I.  Preliminary  examination. 

II.  Distil  in  the  presence  of  an  acid  with  the  aid  of  steam 
the  volatile  poisons,  hydrocyanic  acid,  phenol  (carbolic  acid) , 
phosphorus,  chloroform,  chloral  hydrate. 


208  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

III.  Extract  the  residue  with  alcohol. 

1.  The  alkaloids. 

2.  Oxalic  and  meconic  acids  and  some  mercury  salts. 

IV.  Destroy  the  organic  matter  and  test  the  solution  for 
metals  of  which  the  compounds  are  poisonous. 

I.  Preliminary  Examination 

Before  proceeding  to  the  systematic  analysis  a  preliminary 
examination  will  be  advantageous.  This  may  be  both  phy- 
sical and  chemical.  In  the  former,  if  the  material  is  a  com- 
plex mixture,  like  the  contents  of  a  stomach  or  vomited 
matter,  its  appearance  should  be  noted,  as  to  whether  there  is 
evidence  of  the  presence  of  unusual  substances.  Sometimes 
the  unaltered  poison  (vegetable  or  animal)  may  be  thus  de- 
tected. A  low-power  magnifying  glass  is  of  service.  Phos- 
phorus will  be  visible  if  the  room  is  darkened,  especially  after 
the  substance  'is  dried  and  stirred  in  the  air.  Any  odors 
should  be  noticed,  as  they  may  be  characteristic  of  some  com- 
pound sought.  Not  more  than  one-tenth  of  the  available 
material  should  be  destroyed  in  preliminary  tests. 

Reaction. — A  strong  acid  or  alkaline  reaction  to  litmus 
paper  may  be  due  to  mineral  or  oxalic  acids  or  to  a  caustic 
alkali. 

The  above  acids  will  be  found  in  the  aqueous  extract  and 
change  a  Congo-red  solution  to  a  deep  blue.  Oxalic  acid 
changes  the  color  of  a  methyl-violet  solution  to  a  blue.  Hy- 
drochloric acid  of  the  concentration  found  in  normal  gastric 
juice  will  do  the  same,  also  very  dilute  solutions  of  nitric  or 
sulphuric  acids.  In  any  considerable  amount  the  last  three 
give  with  methyl- violet  a  green  to  yellow  color.  If  a  few 
cubic  centimeters  of  the  solution  be  slowly  evaporated  from  a 
small  test-tube  with  a  few  crystals  of  cane-sugar  there  will  be 
with  sulphuric  acid  a  blackening  or  charring  as  the  water  is 
driven  off;  with  nitric,  an  evolution  of  yellowish-brown  oxids 


THE   DETECTION    OF   POISONS  2OQ 

of  nitrogen;  with  hydrochloric,  a  yellow,  then  brown  color  in 
the  liquid,  with  the  odor  of  the  acid  gas;  with  oxalic  acid,  no 
odor  or  colored  gas,  nor  discoloration  until  the  water  has  been 
expelled,  when  the  sugar  will  be  converted  to  caramel  if  the 
heating  is  continued. 

Of  the  caustic  alkalies  ammonium  hydroxid  will  have  already 
been  recognized  by  its  odor  although  it  must  be  remembered 
that  ammonia  may  be  formed  by  the  bacterial  decomposi- 
tion of  nitrogenous  matter.  The  sodium  hydroxid  and 
potassium  hydroxid  can  be  dissolved  from  dried  residues  by 
hot  alcohol.  They  remain  after  this  has  been  distilled. 
Solutions  of  these  turn  litmus-paper  blue  and  give  a  red  color 
to  a  phenolphthalein  solution.  The  alkaline  carbonates  do 
the  same,  but  not  after  an  excess  of  barium  chlorid  has  been 
added  to  the  solution.  The  alkaline  hydroxids  will  act  upon 
litmus  and  phenolphthalein  after  the  addition  of  barium 
chlorid  as  well  as  before. 

Test  for  phosphorus  and  hydrocyanic  acid  in  a  partly  filled 
flask  with  lead  acetate  paper  (437)  and  silver  nitrate  paper 
(437),  also  guaiacum  paper  (432)  and  watch-glass  tests  (433 
and  436). 

Dialysis  of  a  portion  of  the  mass  may  render  possible  the 
classification  of  the  poison.  Place  the  substance  in  a  dialyzer, l 
after  thinning  with  water  and  slightly  acidifying  with  nitric 
acid,  and  let  this  stand  24  hours  in  four  or  five  times  its  vol- 
ume of  water.  The  liquid  outside  is  then  concentrated,  if  de- 
sirable, and  examined  for  some  of  the  groups  of  poisons,  one 
portion  being  tested  by  hydrogen  sulphid  for  the  heavy 
metals,  another  for  the  alkaloids,  by  means  of  the  general 
reagents  (page  139),  another  for  the  acids  of  alkalies,  etc. 

1  The  dialyzer  may  consist  of  a  wide  glass  tube  which  has  one  end  closed 
with  a  membrane  of  animal  or  vegetable  parchment.  The  substance  to  be 
dialyzed  is  placed  in  this  and  the  whole  is  suspended  in  a  vessel  of  water  so 
that  the  level  inside  and  out  is  the  same.  Or,  instead  of  this,  a  parchment 
tube  containing  the  substance  may  be  suspended  from  the  ends  in  the  water. 

14 


2IO 


INTORDUCTION   TO   CHEMICAL   ANALYSIS 


If  in  the  preliminary  tests  some  indication  has  been  ob- 
tained of  the  class  of  compounds  the  search  should  be  made 
first  for  these.  If  no  indication  had  been  obtained  of  the  class 
to  which  the  poison  may  belong  a  systematic  examination 
must  be  undertaken  for  all  that  can  be  present.  For  this  the 
remaining  material  may  be  divided  into  four  parts  and  one 


FIG.  18. — Two  forms  of  dialyzers;  the  first  an  open  glass  tube  with  parch- 
ment end,  the  second  a  tube  of  parchment  paper. 

examined  for  each  class  of  compounds  unless  the  poison  is 
found  sooner,  one  being  reserved  for  emergencies  or  con- 
firmatory tests.  The  plan  of  analysis  given  will  serve  as  a 
general  guide  for  work,  although  the  chemist  will  of  ten  modify 
it  in  accordance  with  his  observations  during  the  progress  of 
the  analysis.  It  is  far  from  including  all  poisons,  but  it  may 
be  used  in  the  search  for  the  more  common.  For  a  complete 


THE    DETECTION   OF   POISONS  211 

course  of  analysis  reference  should  be  made  to  some  larger 
work,  such  as  that  of  Haines  and  Petersen  or  Dragendorff. 

II.  Poisons  Distilled  from  an  Acid  Liquid 

Dilute  the  finely  divided  mixture  into  a  thin  fluid  with  water, 
then  add  enough  dilute  sulphuric  acid  to  merely  give  an  acid 
reaction.  Distil  from  a  retort  on  a  water-bath,  toward  the 
last  of  the  operation  passing  in  steam  from  a  flask  through  a 
tube  which  extends  nearly  to  the  bottom  of  the  liquid.  The 
distillate  must  be  well  cooled  by  a  long  Liebig's  condenser. 

Hydrocyanic  Acid 

(Ferrocyanids  and  ferricyanids  must  be  absent,  or  the  tests 
are  not  conclusive). 

The  volatile  acid  has  the  characteristic  odor  of  bitter 
almonds,  and  appears  in  the  first  stages  of  the  distillation. 

432.  The  gas  or  a  drop  of  the  distillate  turns  blue  a  white 
paper  previously  moistened  with  very  dilute  copper  sulphate 
solution  and  an  alcoholic  solution  of  guaiacum  resin.     There 
are  a  few  other  substances  which  give  a  similar  result,  but 
they  are  not  commonly  present. 

433.  With  silver  nitrate  a  white  precipitate  is  produced, 
which  dissolves  in  ammonium  hydroxid  or  sodium  hydroxid 
and  is  reprecipitated  therefrom  by  nitric  acid  in  which  it  is 
insoluble.     It  is  not  discolored  by  light  like  silver  chlorid 
(225).     To  remove  any  hydrochloric  acid  which  may  have 
been  distilled  over  from  the  retort,  before  testing  with  the 
silver  nitrate,  the  distillate  can  be  shaken  with  a  little  pow- 
dered borax  and  redistilled.     For  very  minute  quantities 
spread  a  drop  of  the  silver  nitrate  solution  over  the  convex 
surface  of  a  watch-glass,  letting  this  stand  over  a  similar  glass 
containing  the  hydrocyanic  acid. 

434.  The  distillate  forms  Prussian  blue  with  iron  salts 
when  treated  according  to  233. 


212 


INTRODUCTION   TO    CHEMICAL   ANALYSIS 


435.  To  a  few  drops  of  the  distillate  add  2  or  3  drops  of 
a  solution  of  potassium  nitrite  and  as  much  ferric  chlorid. 
Drop  in  dilute  sulphuric  acid  until  the  brown  color  becomes 


FIG.  19. — Apparatus  for  distillation  with  the  aid  of  steam.  This  is  passed  into 
the  liquid  in  the  retort  by  a  tube  which  enters  through  the  tubulure.  The 
flask  at  the  right  in  which  the  steam  is  generated  is  provided  with  a  straight 
safety-tube,  open  at  both  ends.  This  allows  air  to  enter  if  the  heating  is 
interrupted  and  thus  prevents  the  liquid  in  the  retort  from  being  drawn  back. 
The  neck  of  the  retort  is  connected  with  a  Liebig's  condenser  and  the  whole 
apparatus  stands  in  a  hood  which  is  provided  with  a  flue  and  a  sliding  door  so 
that  obnoxious  gases  are  removed  from  the  room. 

a  yellow  and  warm  gently.     Cool,  precipitate  the  iron  with 
a  few  drops  of  ammonium  hydroxid,  filter,  and  to  the  filtrate 


THE   DETECTION    OF   POISONS  213 

add  one  drop  of  very  dilute  ammonium  sulphid.  A  violet 
color  appears,  changing  to  blue,  green,  and  yellow.  The 
hydrocyanic  acid  has  been  converted  into  a  nitroprussid 
which  gives  the  above  results  with  sulphids.  Less  than  one 
part  in  a  million  can  be  thus  detected. 

436.  Place  between  two  watch-glasses  as  in  433  adding 
H2SO4  for  a  cyanid.  Moisten  the  under  surface  of  the  upper 
glass  with  yellow  ammonium  sulphid  solution  and  allow  to 
stand  10  minutes.  Ammonium  sulphocyanate  is  formed  on 
the  upper  glass.  Remove  this  and  dry  on  a  water-bath. 
Very  dilute  ferric  chlorid  gives  a  red  color  with  the  residue 
after  acidifying  with  HC1  (57). 

Phenol  (Carbolic  Acid) 

As  this  is  less  volatile  than  the  hydrocyanic  acid,  the  latter 
portion  of  the  distillate  can  be  used  for  testing.  With  this  try 
the  phenol  reactions  (page  127). 

Chloroform 

If  much  is  present  it  will  appear  as  colorless  drops  with 
the  characteristic  odor,  unless  the  latter  is  hidden  by  other 
compounds.  Try  the  test  with  anilin  and  potassium  hydroxid 
(259)  and,  if  necessary,  confirm  by  260. 

Chloral  Hydrate 

The  distillate  gives  no  odor  of  chloroform  until  it  has  been 
warmed  with  an  alkali.  Test  with  resorcin  (301),  ammonium 
sulphid  (302),  or  with  sodium  hydroxid  (298). 

Phosphorus 

Carry  on  the  distillation  in  a  perfectly  dark  room.  Notice 
the  luminous  ring  in  the  condenser  (438).  For  a  confirma- 
tory test  use  the  color  of  the  flame  with  hydrogen  generator 
(439),  also  effect  on  silver  nitrate  paper  (437)  as  directed  below. 


214  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

The  yellow  variety  of  phosphorus  is  the  one  which  acts 
poisonously.  Its  odor  and  luminous  appearance  in  the  dark 
are  characteristic. 

437.  When  an  acidified  liquid  or  mass  containing  phos- 
phorus is  placed  in  a  flask  of  such  a  size  that  it  shall  be  at  least 
three-fourths  filled  with  air  and  from  the  loosely  inserted 
stopper  strips  of  filter-paper  are  hung  which  have  been  moist- 
ened, one  with,  a  solution  of  lead  acetate  made  alkaline  with 
sodium  hydroxid,  and  the  other  with  silver  nitrate,  the  latter 
is  darkened  slowly  by  the  reduction  of  the  silver  salt  by  the 
vapors  of  phosphorus.     The  former  is  unaffected.     Hydrogen 
sulphid  blackens  both  papers  and  its  presence  destroys  the 
value  of  the  test  unless  it  is  held  in  the  solution  by  the  use  of 
sodium  hydroxid.     Failure  of  the  paper  to  darken  indicates 
the  absence  of  phosphorus.     The  flask  should  be  allowed  to 
stand  a  considerable  time  in  a  dark  place  before  the  conclu- 
sion of  the  test. 

438.  If  a  mixture  containing  phosphorus  is  diluted  with 
water  then,  after  acidifying,  is  placed  in  a  retort  and  distilled 
in  a  dark  room  by  passing  steam  through  it,  the  vapors  being 
cooled  by  a  Liebig's  condenser,  at  the  point  where  the  steam 
is  condensed  there  appears  a  phosphorescent  ring.     This 
moves  back  and  forth  in  the  condenser  and  may  be  visible  a 
long  time.     Very  minute  amounts  of  phosphorus  may  be  thus 
detected,  but  the  phosphorescence  is  prevented  by  mercuric 
chlorid,  alcohol,  and  some  other  volatile  substances.     It  ap- 
pears, however,  when  these  latter  have  been  removed. 

439.  When  treated  in  a  Marsh's  apparatus,  as  described 
in  the  test  for  arsenic  (112),  phosphorus  forms  hydrogen 
phosphid,  PHs,  which,  when  passed  into  silver  nitrate  solution, 
forms  dark,  insoluble  silver  phosphid  mixed  with  metallic 
silver.     This  compound  can  be  filtered  out  and  tested,  after 
acidifying,  by  distillation  as  above.     If  the  gas  is  passed 
through  a  tube  filled  with  pieces  of  pumice  saturated  with 


THE   DETECTION    OF    POISONS  215 

sodium  hydroxid  for  the  purpose  of  absorbing  the  hydrogen 
sulphid  which  may  be  present,  the  flame  of  the  ignited  gas  is 
greenish.  It  should  not  be  lighted  until  the  air  has  been 
completely  expelled  and  should  be  burned  from  a  small 
metal  tube,  for  instance,  a  mouth  blowpipe,  to  avoid  the 
color  which  the  sodium  of  the  glass  would  impart.  The  color 
can  be  made  more  perceptible  by  holding  a  cold  porcelain 
dish  in  the  flame. 

HI,  i.  Examination  for  Alkaloids 

To  another  portion  of  the  finely  divided  mass  add  so  much 
dilute  sulphuric  acid  that  the  reaction  is  faintly  acid  (but  not 
more  than  5  c.c.  of  20  per  cent,  acid  for  each  100  c.c.  of  the 
liquid) .  Mix  thoroughly  and  let  it  stand  several  hours  at  40° 
to  50°,  then  filter  throughacloth.  Repeat  the  extraction,  unit- 
ing the  filtrates.  Evaporate  on  a  steam-bath  to  a  thin  syrup, 
but  no  farther;  dilute  with  3-4  times  its  volume  of  90-95  per 
cent,  alcohol,  which  dissolves  the  alkaloidal  sulphates,  and 
filter  after  it  has  stood  24  hours.  Pour  the  filtrate  into  a 
flask  and  distil  off  the  alcohol  on  a  steam-bath.  The  sulphates 
of  the  alkaloids  remain.  Test  a  portion  for  these,  reserving 
some  for  III,  2;  III,  3;  and  III,  4. 

The  alkaloids  and  some  other  organic  poisons  can  be,  to 
some  extent,  classified  and  separated  by  the  differences  in 
solubility  of  their  salts  and  the  free  substance  in  certain  organic 
solvents.  When  the  aqueous  liquid  is  mixed,  without  too 
violent  shaking,  with  the  solvent  and  the  mixture  is  allowed  to 
stand,  the  solvent,  which  is  selected  so  as  to  be  lighter  or  heavier 
than  the  original  fluid,  removes  from  this  liquid  the  alkaloid. 
B  y  performing  this  operation  in  a  separatory  funnel  or  a  burette 
with  a  glass  stop-cock  one  liquid  can  be  separated  from 
the  other.  To  ascertain  if  anything  has  dissolved,  or  to 
test  it,  a  drop  of  the  solvent  after  separation  can  be 


2l6 


INTRODUCTION    TO    CHEMICAL   ANALYSIS 


allowed  to  evaporate  on  a  watch  crystal.  The  steps  in  this 
separation  are: 

|b,  (a)  If  necessary  add  water  until  the  acidified  solution  is  not 
syrupy,  then  shake  in  a  separatory  funnel  with  petroleum 
ether,  repeating  until  the  soluble  matter  has  been  dissolved. 

The  petroleum  ether  contains 
some  of  the  coloring  matters,  also 
the  fats,  and  the  phenol  which 
was  not  removed  by  distillation. 
(b)  To  the  solution  from  which 
the  petroleum  ether  has  been 
separated  add  sufficient  ammo- 
nium hydroxid  to  make  it 
strongly  alkaline.  The  alkaloids 
are  set  free  and,  if  in  large 
amounts,  may  appear  as  a  precipi- 
tate. Whether  this  is  so  or  not, 
shake  the  liquid  with  more  petro- 
leum ether,  separating  as  before, 
after  standing  long  enough  for  the 
petroleum  ether  to  have  com- 
pletely risen  to  the  top  of  the 
aqueous  solution.  If  a  residue 
remains  when  the  petroleum  ether 
has  spontaneously  evaporated  in 
a  watch-glass,  it  may  be  cocaine 
which  is  in  white  crystals.  If 

,  .,  .  ,-1 

such  a  residue  remains  treat  the 

\{QU{^  with   the   petroleum  ether 

as  long  as  a  drop  leaves  any  solid 
when  evaporated.  Test  the  residues  by  the  cocain  reactions 
(page  144). 

(c)  In  the  same  manner  shake  the  remaining  solution  from 
the  petroleum  ether  extraction  with  benzene  (benzol,  C6H6) 


FIG.     20.  —  Different    forrrs    of 

separatory  funne'.s  used  for  sep- 

arating  a  liquid  from  another  in 
which  it  is  insoluble. 


THE    DETECTION    OF    POISONS 


217 


letting  the  benzene  evaporate  on  a  number  of  watch-glasses 
without  heating.  To  learn  if  an  alkaloid  is  present  dissolve 
in  water  with  a  drop  of  acid  and  apply  the  general  alkaloidal 
reagents  (page  139). 

There  may  be  contained  in  the  residues  from  the  benzene 
atropin,  brucin,  narcotin,  strychnin,  and  veratrin.  Observe 
the  appearance  and,  if  the  general  alkaloidal  reagents  (page 
139)  show  the  presence  of  alkaloids,  test  the  separate  portions 
with  a  drop  of  concentrated  sulphuric  acid  and  another  of 
concentrated  nitric  acid. 


Usual  Form. 

H2SO3, 
Concentrated. 

HN03, 

Concentrated. 

Tests 
on  Page. 

Atropin, 

crystalline, 

colorless, 

colorless, 

143 

Brucin, 

amorphous, 

colorless. 

deep  red, 

142 

Narcotin, 

crystalline, 

colorless, 

yellow, 

141 

Strychnin, 

crystalline, 

colorless, 

yellow, 

141 

Veratrin, 

amorphous, 

yellow,       orange, 
then  red  (396) 

yellow  to  colorless. 

144 

Confirm  the  indicated  alkaloids  by  the  individual  reactions. 

(d)  The  solution  which  has  been  treated  with  benzene  is  to 
be  shaken  with  amyl  alcohol  and,  after  separating,  the  latter 
is  to  be  evaporated  on  the  steam-bath. 

Morphin  remains  as  a  crystalline  residue.  Confirm  it  by 
its  characteristic  reactions  (page  140). 

HI,  2.  Examination  for  Oxalic  Acid 

A  part  of  the  alcohol  solution  from  the  dried  substance 
is  filtered  and  evaporated  to  dryness  on  the  steam-bath  and 
the  residue  dissolved  in  water.  This  solution,  if  oxalic  acid 
is  present,  gives  calcium  oxalate  with  calcium  chlorid  (176); 
it  decolorizes  potassium  permanganate  (177);  oxalic  acid 
precipitates  from  a  solution  of  gold  chlorid  metallic  gold  in 
scales  (133).  Other  tests  have  been  referred  to  in  the  pre- 
liminary examination. 


2l8  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

ffl,  3.  Examination  for  Meconic  Acid 

Test  with  silver  nitrate  (315)  and  ferric  chlorid  (314). 
The  latter  is  the  more  characteristic. 

in,  4.  Examination  for  Metallic   Compounds   Soluble  in 
Alcohol,  Especially  those  of  Mercury 

Although  the  general  tests  for  the  metallic  compounds 
are  made  with  another  portion  of  the  substance,  this  solution 
may  be  tested  with  hydrogen  sulphid  and  the  test  completed 
in  the  usual  manner  (Table  III).  Since  only  substances  solu- 
ble in  alcohol  can  be  present  here  the  discovery  of  a  metal 
affords  some  indication  of  the  original  compound. 

IV.  Metallic  Poisons 

They  may  include  compounds  of  arsenic,  antimony,  lead, 
copper,  mercury,  silver,  zinc,  and  barium. 

These  may  be  sought  in  the  residue  from  the  alcoholic 
extraction  of  III  (except  lead  and  barium)  or  another  portion 
of  the  mass  under  investigation.  In  order  to  obtain  them  in 
a  sufficiently  pure  state  to  identify  them  any  organic  matter 
present  should  be  first  destroyed.  This  is  accomplished 
through  oxidation  by  means  of  chlorin.  The  finely  divided 
mass  is  diluted  to  a  thin  mixture  with  pure,  concentrated 
hydrochloric  acid,  and  a  little  water.  Place  it  in  a  flask 
which  it  must  no  more  than  half  fill,  and  drop  in  from  time 
to  time  a  little  powdered  potassium  chlorate  as  long  as  the 
fluid  has  a  dark  color,  assisting  the  oxidation  toward  the  last 
by  warming  on  the  steam-bath.  Avoid  the  use  of  unneces- 
sary amounts  of  either  acid  or  chlorate.  Cellulose  or  masses 
of  fat,  which  are  difficult  to  decompose,  may  be  filtered  out 
and  the  filtrate  employed  for  the  examination. 

The  clear  liquid  is  freed  from  chlorin  by  passing  carbon 


THE   DETECTION    OF,  POISONS  2IQ 

dioxid  through  it  and  afterward  saturated  with  hydrogen  sui- 
phid.  This  may  be  done  in  a  flask  fitted  with  a  two-holed 
rubber  stopper  and  two  tubes.  The  entrance  tube  is  con- 
nected with  the  generator  and  extends  to  the  bottom  of  the 
solution.  Gas  is  passed  in  for  24  hours  (102) ,  the  liquid  being 
warmed  and  shaken  occasionally.  The  zinc  and  barium  alone 
remain  unprecipitated.  Filter  and  wash. 

Compare  the  color  of  the  precipitate  with  that  of  the  sul- 
phids  of  metals  possibly  present.  Separate  or  test  it  by  the 
methods  given  in  Table  IV,  confirming  by  the  reaction  of  the 
individual  metals,  using  for  the  latter  purpose  a  fresh  portion 
of  the  solution  if  desirable. 

Test  the  filtrate  for  zinc  after  the  method  of  Table  II  and 
the  zinc  reactions  (page  39).  If  this  is  not  present  or  has 
been  removed,  test  for  barium  according  to  Table  I  and  the 
barium  reactions  (page  30). 

Special  Problems  in  Toxicology 

Detection  of  Arsenic  in  Milk  or  Liquid  Foods. — Use 
Reinsch's  test  (108). 

Detection  of  Arsenic  or  Lead  in  Urine. — Acidify  about  a  liter 
with  hydrochloric  acid,  saturate  with  hydrogen  sulphid  gas 
and  let  it  stand  in  a  corked  flask  for  24  hours.  If  there  is  a 
precipitate,  filter  and  wash  it.  If  it  is  yellowish,  test  for 
arsenic  by  Gutzeit's  test  (109).  If  black  dissolve  in  a  few 
drops  of  nitric  acid  and  test  for  lead  (147,  etc.). 

Detection  of  Mercury  in  Urine. — Acidify  about  a  liter  of 
urine  with  i  or  2  c.c.  of  hydrochloric  acid,  warm  to  5o°-6o°, 
add  half  a  gramme  of  zinc  dust  or  freshly  precipitated  metal- 
lic copper  (74).  Stir  vigorously  half  a  minute  and  then  filter. 
Metallic  mercury  is  precipitated  (.88).  .Wash  and  confirm 
by  heating  in  a  tube  (88) . 

Detection  of  Copper  in  Coloring  Matter. — Digest  the  fabric 


220  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

some  time  with  concentrated  ammonium  hydroxid.  A  blue 
solution  indicates  copper  (71). 

Detection  of  Copper  in  Foods. — A  large  amount  will  be  pre- 
cipitated from  the  mass,  after  acidifying  with  hydrochloric 
acid,  by  means  of  iron,  like  a  bright  knife-blade  (74).  When 
washed  and  digested  with  concentrated  ammonium  hydroxid 
it  makes  a  blue  solution.  If  the  amount  of  copper  is  small, 
it  may  be  necessary  to  first  destroy  the  organic  matter  by 
potassium  chlorate  (page  218). 

Detection  of  Lead  in  Water. — Concentrate  a  large  volume  of 
water  (one  or  more  liters)  to  100  c.c.  in  a  dish  with  lead-free 
glazing,  after  adding  one  drop  of  nitric  for  each  liter.  If 
it  is  discolored  by  organic  matter,  remove  this  by  potassium 
chlorate  and  hydrochloric  acid  (page  218).  Then  saturate 
with  hydrogen  sulphid  gas,  heat  to  boiling,  filter  and,  after  dis- 
solving in  nitric  acid,  confirm  by  lead  tests  (147,  etc.). 

Detection  of  Dangerous  Amounts  of  Lead  in  Glazing  or 
Enamel. — In  the  vessel  to  be  tested  boil  for  half  an  hour  4  per 
cent,  acetic  acid,  adding  water  as  it  evaporates.  Test  the 
liquid  after  this  time  with  hydrogen  sulphid  as  above. 


ANALYSIS   BY   MEANS    OF   THE   BLOWPIPE  221 


CHAPTER  III 

ANALYSIS  BY  MEANS  OF  THE  BLOWPIPE 

MANY  of  the  important  physical  and  chemical  properties  of 
the  metals  and  their  alloys  which  have  not  been  shown  by  the 
preceding  work  can  be  demonstrated  by  the  use  of  the  blow- 
pipe (page  15).  These  are  often  of  the  greatest  importance 
in  the  technical  use  of  the  metals  or  the  manufacture  of  their 
alloys,  and  they  can  best  be  learned  by  careful  observation 
and  long  practice.  They  are  such  as  are  modified  or  are  im- 
parted to  the  metals  by  heating.  By  the  same  means  some  of 
the  non-metallic  elements  and  compounds  undergo  charac- 
teristic changes  also. 

Many  metals,  as  well  as  some  non-metallic  substances, 
give  characteristic  reactions  during  or  after  heating  alone  or 
with  some  chemical  agent.  To  obtain  these  first  excavate  a 
shallow  cavity  near  one  end  of  a  stick  of  charcoal;  place  in  this 
some  of  the  substance  equal  to  half  a  pea  in  size,  and  apply  the 
blowpipe  flame,  using  first  the  oxidizing  flame,  then  the  re- 
ducing (pages  15-16),  and  heating  the  charcoal  underneath 
the  assay  as  much  as  possible.  The  flame  should  be  directed 
toward  the  unoccupied  end  of  the  charcoal  so  as  to  allow  for  a 
possible  deposition  of  a  coating. 

In  some  cases  the  results  are  better  observed  if,  instead 
of  the  charcoal,  the  assay  is  heated  on  a  thin  slab  of  plaster 
of  Paris,  made  by  pouring  a  thick  mixture  of  this  with  water 
into  a  pasteboard  box  and  letting  it  harden.  For  light- 
colored  coatings  it  can,  previous  to  using,  be  blackened  by 
smoking. 


INTRODUCTION   TO   CHEMICAL   ANALYSIS 

I.  The  Non-metallic  Substances  and  Metals  of  the  Alka- 
lies and  Alkaline  Earths 

Water,  if  present  in  large  amounts  in  the  form  of  water  of 
crystallization,  as  in  alum,  causes  the  substance  to  melt  easily. 
When  it  has  been  evaporated  by  continued  heating  the  anhy- 
drous substance  remains  on  the  charcoal.  If  only  a  small 
quantity  of  moisture  is  present  it  will  be  necessary  to  heat  the 
powdered  solid  in  a  glass  tube  when  the  moisture  appears  on 
the  cool  walls  of  the  tube.  Most  hydroxids  give  up  their 
water  in  this  way. 

Many  nitrates  and  chlorates  deflagrate  when  heated  on 
charcoal,  that  is,  appear  to  burn  rapidly,  leaving  little  visible 
residue.  This  is  especially  true  of  salts  of  the  alkalies.  Ni- 
trates, when  mixed  with  acid  potassium  sulphate  and  heated 
in  a  small  glass  tube,  evolve  a  yellowish  gas,  best  seen  by 
looking  down  into  the  tube.  This  gas  turns  brown  a  slip  of 
filter-paper  which  has  been  dipped  in  a  solution  of  ferrous 
sulphate. 

Chlorates  yield  oxygen,  sometimes  mixed  with  chlorin, 
when  heated  in  a  tube.  If  cupric  oxid  is  dissolved  in  a  bead 
of  microcosmic  salt  by  the  aid  of  the  oxidizing  flame,  when  a 
chlorate  is  added  the  flame  becomes  azure-blue.  Other 
chlorin  compounds  give  the  above  reaction  with  the  microcosmic 
bead  and  cupric  oxid  as  do  bromin  compounds  also.  Many 
of  them  decrepitate,  that  is,  fly  to  pieces  when  heated  because 
of  the  moisture  that  is  contained  in  the  crystals,  though  this  is 
not  characteristic  of  chlorids. 

Free  sulphur  and  sulpkids  oxidize  on  charcoal  in  the  oxidiz- 
ing flame  with  the  evolution  of  sulphurous  oxid,  recognizable 
by^its  odor. 

Sulphates  and  other  sulphur  compounds,  if  they  are  fused 
on  charcoal  with  dry  sodium  carbonate,  and  if  then  the  fused 
mass  is  placed  on  a  silver  coin  and  moistened,  will  turn  the 


ANALYSIS   BY   MEANS    OF    THE  BLOWPIPE  223 

coin  brown  to  black  from  the  presence  of  sodium  sulphid  in 
the  fused  residue. 

Silicates  when  heated  on  charcoal  give  no  characteristic 
results.  If  fused  in  the  bead  of  microcosmic  salt  their  bases 
dissolve  and  the  silica  floats  in  a  skeleton-like  mass  in  the  hot 
bead. 

Organic  compounds  usually  blacken  when  heated,  although 
they  do  not  always  do  so.  If  they  contain  no  metals  they  are 
completely  combustible  or  volatile. 

Carbon  can  be  burned  completely.  Graphite  is  combustible 
with  difficulty  before  the  blowpipe.  It  deflagrates  when 
heated  with  potassium  chlorate  on  platinum  foil. 

Sodium  and  potassium  compounds  melt  and  sink  into  the 
charcoal.  Sodium  gives  a  yellow  flame;  potassium,  a  violet 
one  (pages  24-25). 

Ammonium  compounds  volatilize  completely,  frequently 
leaving  a  distinct  white  coating  on  the  charcoal. 

The  compounds  of  the  alkaline  earths,  barium,  strontium, 
calcium,  and  magnesium,  remain  white  on  ignition  and  often 
become  highly  luminous.  They  are  sometimes  infusible  and 
sometimes  melt  and  sink  into  the  charcoal.  Compounds  of 
these  metals  that  reduce  to  the  oxids  will  turn  moist  litmus- 
paper  blue  when  placed  upon  it. 

Barium  when  held  in  the  blue  flame  colors  it  yellowish- 
green.  This  may  be  made  more  perceptible  by  moistening 
the  substance  with  a  drop  of  hydrochloric  acid.  In  the  same 
way 

Strontium  gives  a  crimson-red,  and 

Calcium  gives  a  yellowish-red  flame. 

Magnesium,  when  ignited,  then  moistened  with  cobalt 
nitrate  solution  and  again  ignited,  is  colored  a  faint  flesh-pink. 

Tests  useful  with  Alloys  and  Amalgams 

The  student  is  to  try  the  following  tests  and  arrange  the 
properties  of  each  metal  thus  ascertained  in  a  tabular  form. 


224  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

1 .  Mix  as  much  of  the  oxid  or  carbonate  of  the  metal  as  is 
equal  to  half  the  bulk  of  a  pea  with  twice  its  volume  of  dry 
sodium  carbonate  and  powdered  borax.     Heat  in  the  reduc- 
ing flame  on  charcoal  as  directed  above.     What  is  the  action 
of  these  reagents  at  high  temperatures?     Note  whether  the 
compounds  are  reducible  under  these  conditions  and  the  ease 
with  which  this  change  can  be  effected.     If  the  metal  is  ob- 
tained note  its  color  and  malleability. 

Test  in  this  manner,  Hg,  As,  Sb,  Sn,  Bi,  Pb,  Cd,  Mn,  Cr, 
Fe,  Co,  Ni,  Cu,  Zn,  Al,  Ag,  Au. 

2.  Heat  on  the  charcoal  stick,  or  on  a  plaster  slab,  a  small 
fragment  of  the  metal,  not  larger  than  one-tenth  the  size  of  a 
pea.     Note  and  tabulate: 

(a)  Degree  of  fusibility. 

(b)  Ease  of  oxidation. 

(c)  Color  of  coating,  hot  and  cold. 

(d)  Volatility  of  metal  or  of  oxid  as  shown  by  position. 
Test  the  same  metals  as  under  i. 

3.  Place  a  small  amount  of  the  oxid  or  carbonate  on  a 
plaster  slab,  moisten  with  a  drop  of  hydriodic  acid,  then  heat 
with  the  oxidizing  flame.     Hold  the  colored  films  over  am- 
monia; touch  them  with  a  drop  of  ammonium  sulphid. 

In  this  manner  test  Hg,  As,  Sb,  Sn,  Bi,  Pb,  Ag.     On  a 
sooted  slab  try  similarly  Cu,  Cd,  Zn. 

4.  After  heating  the  oxid  on  charcoal  or  plaster  moisten 
with  a  drop  of  cobalt  nitrate  solution,  then  heat  thoroughly; 
test  thus  Sb,  Zn,  Sn,  Al. 

5.  Dissolve  a  minute  particle  of  the  oxid  in  a  borax  bead  by 
heating.     Tabulate  the  colors  of  the  beads,  both  hot  and  cold, 
after  heating  in  the  oxidizing,  then  in  the  reducing  flame. 

Test  so  Mn,  Cr,  Fe,  Co,  Ni,  Cu. 

6.  Moisten  a  particle  of  the  oxid  or  carbonate  with  hydro- 
chloric acid  and  heat  in  the  blue  Bunsen  flame  on  a  platinum 
wire  or  an  asbestos  fiber.     Why  is  the  platinum  wire  less 
suitable  as  a  support? 


ANALYSIS   BY   MEANS    OF   THE   BLOWPIPE  225 

Test  so  Cu,  Pb,  Sb. 

7.  Heat  in  a  small  tube  closed  at  one  end,  Hg,  As,  As2O3. 

II.  The  Common  Heavy  Metals  Including  the  Alloys  and 

Amalgams 

i.  Easily  Volatile;  Form  a  Coating  or  Disappear 

Mercury  is  completely  volatilized,  leaving  at  most  a  dis- 
tant, gray  coating.  On  plaster  this  is  very  distinct.  Some 
salts  of  mercury  form  white  coatings.  If  mixed  with  dry 
sodium  carbonate  and  heated  in  a  tube,  globules  of  mercury 
are  deposited  in  the  cool  part  of  the  tube.  On  plaster  with 
sulphur  and  potassium  iodid,  or  HI,  it  gives  a  scarlet  coating 
with  yellow;  if  quickly  heated  this  is  dull  and  black. 

Arsenic  volatilizes  completely  with  distant  white  incrusta- 
tion. It  gives,  when  vaporized,  a  garlic  odor.  In  a  glass 
tube  with  sodium  carbonate,  most  compounds  give  a  dark 
ring  with  a  metallic  luster.  This  is  metallic  arsenic.  No 
globules  can  be  perceived. 

Antimony  gives  brittle,  white  metallic  globules.  These 
produce  white  fumes  when  heated  in  the  oxidizing  flame 
and  a  distinct  white  incrustation  with  a  bluish  border.  The 
metal  continues  to  burn  after  removing  from  the  flame,  finally 
surrounding  itself  with  crystals  of  the  oxid.  If  dropped  on 
the  floor  when  melted  it  rolls  and  leaves  a  white  track  behind. 
It  imparts  a  greenish-blue  color  to  the  oxidizing  flame. 
Heated  on  plaster  with  sulphur  and  potassium  iodid,  or  HI, 
it  gives  an  orange  coating  stippled  with  peach-red. 

Bismuth  gives  reddish- white,  brittle  metallic  globules  with 
a  yellow  coating.  When  mixed  with  sulphur  and  potassium 
iodid,  or  HI,  and  heated  with  a  small,  oxidizing  flame  a  bril- 
liant scarlet  coating  is  produced,  with  a  yellow  one  at  a 
greater  distance;  with  the  last  reagents  on  plaster,  a  choco- 
late-brown coating  appears  with  underlying  scarlet;  with 
ammonia  this  becomes  orange-yellow  and  cherry-red. 

15 


226  INTRODUCTION   TO   CHEMICAL   ANALYSIS 

Lead  forms  white,  malleable  metallic  globules  with  a  yellow 
coating  which  has  a  white  edge.  The  flame  is  a  bluish-white. 
Heated  with  sulphur  and  potassium  iodid,  or  HI,  the  volatile 
coating  is  a  bright  yellow. 

Cadmium  has  a  reddish-brown  incrustation  near  the  assay 
with  a  yellow  one  farther  away.  No  globules  can  be  ob- 
served. If  heated  strongly  in  a  glass  tube  the  metal  volatil- 
izes, condensing  to  bright  globules  on  cooling. 

2.  The  Non-volatile  Metals  or  those  which  Vaporize  with  Difficulty, 
and  which  Color  the  Borax  Bead 

Many  of  these  impart  a  definite  color  to  a  bead  of  fused 
borax,  held  in  a  loop  of  platinum  wire.  To  as  far  as  possible 
avoid  the  presence  of  fusible  metals  which  would  alloy  the 
platinum  the  test  substance  may  be  roasted  on  charcoal  to 
vaporize  the  volatile  ones  and  to  oxidize  most  of  the  others, 
with  any  sulphur  that  may  be  present.  This  is  accomplished 
by  heating  a  thin  layer  with  a  gentle  oxidizing  flame  so  as  not 
to  fuse  it,  turning  the  particles  in  order  to  thoroughly  oxidize 
or  volatilize  them.  Some  of  the  metals  which  may  remain 
cannot  be  oxidized  at  all,  and  in  many  cases  the  oxidation  is 
incomplete. 

At  the  end  of  a  platinum  wire  make  a  loop  about  an  eighth 
of  an  inch  in  diameter  by  winding  it  around  a  match  or  the 
point  of  a  pencil.  Heat  it,  and  after  dipping  it  into  powdered 
borax  [heat  the  mass  until  it  fuses  to  a  clear,  colorless  bead. 
If  this^is  colored,  the  wire  was  not  clean  and  the  bead  should  be 
removed  by  straightening  the  wire  and  a  new  one  made. 
Touch  the  hot  bead  to  a  minute  fragment  of  the  oxidized 
residue  which  will  adhere,  then  heat  again,  first  in  the  oxidiz- 
ing flame,  then  in  the  reducing  flame  of  the  blowpipe.  No- 
tice the  color  which  is  given  to  the  bead.  The  latter  should, 
with  most  metals,  remain  transparent.  If  it  becomes  black 
it  is  because  too  much  of  the  substance  has  been  added. 

Manganese  in  the  oxidizing  flame  gives  a  reddish-purple 


ANALYSIS  BY  MEANS    OF   THE  BLOWPIPE  227 

which  becomes  colorless  in  the  reducing  flame.  A  small  por- 
tion fused  into  a  bead  of  potassium  nitrate  and  sodium  car- 
bonate becomes  deep  green;  a  large  amount  gives  a  black 
color. 

Chromium  gives  an  emerald-green  bead,  there  being  little 
change  produced  by  the  different  flames.  In  the  bead  of 
potassium  nitrate  and  sodium  carbonate,  a  yellow  color  is 
produced.  If  this  mass  is  dissolved  in  a  drop  of  water  it 
makes  a  bright  yellow  solution. 

Cobalt  produces  a  deep  blue  bead,  the  color  being  un- 
changed in  the  reducing  flame.  Heated  alone  on  charcoal  in 
the  reducing  flame  the  metal  is  obtained — a  black  magnetic 
powder. 

Iron  colors  the  bead  yellow  to  reddish  in  the  oxidizing  flame, 
the  shade  being  darker  when  hot,  and  also  as  the  amount  of 
iron  is  increased.  In  the  reducing  flame  it  is  colorless  to 
bottle-green.  Heated  alone  on  charcoal,  iron  and  its  com- 
pounds give  a  black  powder  which  is  attracted  by  a  magnet. 

Nickel  when  oxidized  gives  a  bead,  violet  when  hot  and 
brown  when  cold.  With  the  reducing  flame  it  is  grayish  or 
colorless.  On  the  charcoal  the  reducing  flame  converts  its 
compounds  to  metallic  nickel,  a  black,  magnetic  powder. 
This  shows  a  metallic  luster  on  rubbing  in  a  mortar. 

Copper  colors  the  borax  green  as  it  is  taken  from  the  oxidiz- 
ing flame,  changing  to  a  blue  as  it  cools.  In  the  reducing 
flame  the  bead  is  colorless  unless  a  large  amount  of  the  metal 
is  present,  when  it  becomes  a  brownish-red  and  opaque  on 
cooling.  On  charcoal  compounds  of  copper  are  reduced  to  the 
metal  which  fuses  to  a  malleable  globule.  In  the  reducing 
flame  the  color  of  the  metal  is  seen,  but  on  removing  it  from 
the  flame,  or  on  heating  in  the  oxidizing  flame,  it  is  covered 
with  a  black  oxid.  Copper  gives  a  green  color  to  the  flame; 
the  chlorid,  which  may  be  formed  by  moistening  with  hydro- 
chloric acid,  gives  a  blue  one. 


228  .    INTRODUCTION   TO    CHEMICAL   ANALYSIS 

3.  Metals  which  Volatilize  with  Difficulty  or  not  at  all  and  do  not 
Color  the  Borax  Bead 

After  heating,  in  order  to  as  far  as  possible  remove  the 
volatile  metals,  if  such  are  present,  place  the  residue  on  a 
clean  piece  of  charcoal  and  heat  again  with  the  oxidizing  and 
then  the  reducing  flame. 

Zinc  gives,  near  the  substance,  a  coating  which  is  yellow 
when  hot  and  white  after  it  has  cooled.  Moisten  this  with 
cobalt  nitrate  solution  and  heat  again  to  as  high  a  tempera- 
ture as  can  be  produced.  The  coating  is  changed  to  a  bright 
green. 

Tin  has  a  coating,  on  or  near  the  assay  but  no  metallic 
globules.  It  is  yellow  when  hot  and  white  when  cold. 
Moistened  with  cobalt  nitrate  solution  and  ignited,  it  be- 
comes bluish-green.  On  plaster  with  sulphur  and  potassium 
iodid  it  gives  a  brownish-orange  coating. 

Aluminum  gives  a  non-volatile  incrustation  on  the  sub- 
stance heated,  which  is  white  both  hot  and  cold.  Ignition 
after  moistening  with  cobalt  nitrate  solution  produces  a 
bright  blue  color. 

Silver  gives  white,  malleable  metallic  globules.  A  dark  red 
coating  is  formed  near  the  assay  but  only  after  very  long 
heating. 

Gold  fuses  to  yellow  metallic  globules  without  a  coating. 
It  is  soluble  in  aqua  regia  but  not  in  nitric  acid. 

Platinum  remains  infusible  on  the  charcoal  and  gives  no 
.  coating.  If  reduced  from  its  compounds  it  is  a  fine,  black 
powder.  If  it  is  in  larger  masses  it  is  of  a  white  color. 

Practical  Exercises 

By  blowpipe  analysis  determine  the  metals  in  unknown 
compounds  furnished  by  the  instructors,  then  the  names  of 
unknown  single  metals,  finally  the  composition  of  simple 
alloys. 


ANALYSIS   BY   MEANS    OF    THE   BLOWPIPE  22Q 

To  fix  in  mind  the  more  important  properties  of  the  metals 
the  following  questions  should  in  all  cases  be  answered. 

Ease  of  reduction?  degree  of  fusibility?  of  volatility? 
ease  of  oxidation?  difference  at  high  temperatures  and  the 
ordinary  ones?  If  an  oxid  is  formed  what  is  its  color?  is  this 
changed  with  the  temperature?  is  the  metal  malleable  or 
brittle?  hard  or  soft? 

Try  to  classify  the  metals  on  the  basis  of  each  of  these 
properties  into 

1.  Very  readily  changed. 

2.  Changed  without  much  difficulty. 

3.  Changed  with  difficulty. 

4.  No  change  effected. 

Questions  for  Further  Study  on  Blowpipe  Analysis 

For  what  reason  is  the  charcoal  under  the  test  substances 
heated  to  a  red  heat  in  blowpipe  analysis?  What  chemical 
action  takes  place  when  chlorates  and  nitrates  are  heated  on 
charcoal?  To  what  class  of  agents  do  these  compounds  be- 
long? Where  are  they  used  when  this  property  is  of  value? 
Are  there  any  substances  with  which  they  would  be  incompati- 
ble ?  When  is  a  high  heat  necessary  to  produce  this  chemical 
reaction?  What  is  an  acid  salt?  What  is  the  action  of  the 
acid  potassium  sulphate  when  heated  with  a  nitrate?  Why 
does  cupric  oxid  in  the  presence  of  a  chlorate  give  a  blue  flame 
instead  of  the  ordinary  green  one  ?  How  does  moisture  cause 
decrepitation  ?  Why  is  it  necessary  to  heat  sulphur  in  order  to 
obtain  an  odor?  What  is  the  black  compound  formed  on  a 
silver  coin  when  it  is  brought  into  contact  with  moist  sodium 
sulphid  ?  Why  must  water  be  added  before  the  discoloration 
can  be  produced  ?  What  brings  about  the  change  of  a  sulphate 
to  a  sulphid  when  treated  in  this  manner?  What  becomes 
of  the  bases  of  a  silicate  when  it  is  heated  with  microcosmic  salt  ? 


230  INTRODUCTION   TO    CHEMICAL    ANALYSIS 

What  is  the  black  substance  produced  when  organic  com- 
pounds are  heated?  What  compounds  are  formed  when 
graphite  is  deflagrated  with  potassium  chlorate?  When 
ammonium  chlorid  is  volatilized  on  charcoal  is  it  a  complete 
chemical  decomposition  or  a  physical  change?  Is  calcium 
oxid  fusible?  What  application  is  made  of  this  property  and 
the  fact  that  it  is  very  luminous  at  high  temperatures  ?  Why 
does  the  addition  of  hydrochloric  acid  to  some  of  these  com- 
pounds make  a  brighter  flame?  How  do  the  compounds  of 
the  alkaline  earths  compare  with  those  of  the  alkali  metals  in 
alkalinity?  Are  there  any  other  of  their  compounds  except 
the  oxids  which  will  have  the  same  effect  on  litmus-paper? 

What  is  the  action  of  sodium  carbonate  in  setting  mercury 
free  from  its  salts?  Which  would  be  the  better  way  to  detect 
small  quantities  of  mercury  by  heating  on  charcoal  or  in  a 
small  tube?  Is  the  sodium  carbonate  used  for  the  same  pur- 
pose with  arsenic  compounds?  How  would  you  explain  the 
absence  of  metallic  globules  when  cadmium,  zinc,  tin  and 
aluminum  compounds  are  heated  on  charcoal  and  their  pro- 
duction when  the  same  process  is  applied  to  compounds  of 
lead  or  bismuth?  Why  is  platinum  wire  used  for  supporting 
the  borax  bead  in  preference  to  wire  of  other  metals?  What 
causes  the  swelling  of  the  powdered  borax  in  making  the 
transparent  bead?  How  does  the  action  of  the  borax  in  the 
bead  explain  its  usefulness  in  the  soldering  of  two  pieces  of 
metal?  How  can  you  explain  the  fact  that  some  metals, 
like  iron,  give  one  color  when  the  bead  is  heated  in  the  oxidiz- 
ing flame,  and  a  different  one  in  the  reducing  flame?  What 
does  the  position  of  the  coating  on  the  charcoal  show  as 
regards  the  volatility  of  the  metals? 


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240  INTRODUCTION   TO    CHEMICAL   ANALYSIS 

THE  CHEMICAL  ELEMENTS^ 

With  their  Symbols  and  Atomic  Weights 

THESE  are  from  the  report  of  the  International  Committee 
on  Atomic  Weights,  1912. 


NAME. 

SYM- 

BOL. 

Aluminum 

Al 

Antimony 

Sb 

Argon 

A 

Arsenic 

As 

Barium 

Ba 

Beryllium1 

Be 

Bismuth 

Bi 

Boron 

B 

Bromin 

Br 

Cadmium 

Cd 

Caesium 

Cs 

Calcium 

Ca 

Carbon 

C 

Cerium 

Ce 

Chlorin 

Cl 

Chromium 

Cr 

Cobalt 

Co 

Copper 

Cu 

Dysprosium 
Erbium 

% 

Europium 

Eu 

Fluorin 

F 

Gadolinium 

Gd 

Gallium 

Ga 

Germanium 

Ge 

Gold 

Au 

Helium 

He 

Hydrogen 

H 

Indium 

In 

lodin 

I 

Iron 

Fe 

Iridium 

Ir 

Krypton 

Kr 

Lanthanum 

La 

Lead 

Pb 

Lithium 

Li 

Lutecium 

Lu 

Magnesium 

Mg 

Manganese 

Mn 

Mercury 

Hg 

Molybdenum      Mo 


AT.  WT. 

NAME. 

SYM- 

AT. WT 

BOL. 

27.1 

Neodymium 

Nd 

144-3 

I  2O.  2 

Neon 

Ne 

2O.  2 

39-88 

Nickel 

Ni 

58.68 

74.96 

Niobium2 

Nb 

93-5 

137-37 

Nitrogen 

N 

14.01 

9.1 

Osmium 

Os 

190.9 

208.0 

Oxygen 

0 

16.0 

ii  .0 

Palladium 

Pd 

106.  7 

79-92 

Phosphorus 

P 

31.04 

112.4 

Platinum 

Pt 

195.2 

132.81 

Potassium 

K 

39-i 

40.07 

Praseodymium 

Pr 

140.6 

12.0 

Radium 

Ra 

226.4 

140.25 

Rhodium 

Rh 

102.9- 

35.46 

Rubidium 

Rb 

85.45 

52.0 

Ruthenium 

Ru 

101  .  7 

58.97 

Samarium 

Sa 

150.4 

Scandium 

Sc 

44.1 

162.5 

Selenium 

Se 

79-2 

167.7 

Silicon 

Si 

28.3 

152.0 

Silver 

Ag 

107.88 

19.0 

Sodium 

Na 

23.0 

157-3 

Strontium 

Sr 

87.63 

69.9 

Sulphur 

S 

32.07 

72-5 

Tantalum 

Ta 

181.5 

197.2 

Tellurium 

Te 

127.5 

3-99 

Thallium 

Tl 

204.0 

1.008 

Thorium 

Th 

232.4 

114.8 

Thulium 

Tu 

168.5 

126.92 

Tin 

Sn 

119.0 

55.85 

Titanium 

Ti 

48.1 

Tungsten 

W 

184.0 

82  .92 

Uranium 

U 

238-5 

139.0 

Vanadium 

V 

51.2 

207.  i 

Xenon 

X 

130.2 

6.94 

Ytterbium 

Yb 

172.0 

174.0 

(Neo-ytterbium) 

24.32 

Yttrium 

Y 

89.0 

54-93 

Zinc 

Zn 

65.37 

200.  6 

Zirconium 

Zr 

90.6 

96.0 

1  Also  called  glucinum. 

2  Also  called  columbium. 


THE   METRIC   SYSTEM  241 

The  Metric  System 

IN  all  work  in  chemistry  the  metric  system  of  weights  and 
measures  is  used  instead  of  the  older  ones  formerly  in  vogue. 
It  is  based  upon  the  meter  (39.37  inches)  which  is  the  unit  of 
length.  The  hundredth  part  of  this  is  the  centimeter  and  a 
cube  of  water  at  4°  C.,  each  side  of  which  is  a  centimeter, 
weighs  a  gramme — the  unit  of  weight.  One-tenth  the  length 
of  the  meter  is  a  decimeter  and  a  cube  which  has  a  side  of  this 
dimension  contains  a  liter — the  unit  of  capacity.  The  frac- 
tions and  multiples  of  these  are  the  following: 


Measures  of  Length 

10  millimeters  (mm).  =  i  centimeter  (cm.). 

10  centimeters  =  i  decimeter. 

10  decimeters  =  i  meter  (m.). 

10  meters  =  i  decameter. 

10  decameters  =  i  hectometer. 

10  hectometers  =  i  kilometer. 


Measures  of  Weight 

10  milligrammes  (mg.)  =  i  centigramme. 

10  centigrammes  =  i  decigramme. 

10  decigrammes  =  i  gramme  (g.  or  gm.). 

10  grammes1  =  i  decagramme. 

10  decagrammes  =  i  hectogramme. 

10  hectogrammes  =  i  kilogramme  (kilo.). 

1  The  me  is  often  dropped  in  the  names  of  these  denominations.     It  is  per- 
haps preferable  that  it  should  be  retained  to  avoid  any  danger  of  reading 
grain  for  gram. 
16 


242  INTRODUCTION   TO    CHEMICAL    ANALYSIS 

Measures  of  Volume 

10  milliliters   =  i  centiliter. 
10  centiliters  =  i  deciliter. 
10  deciliters    =  i  liter  (1.). 
10  liters  =  i  decaliter. 

10  decaliters   =  i  hectoliter. 
10  hectoliters  =  i  kiloliter. 

In  volumetric  analysis  it  is  important  to  remember  that 
one  cubic  centimeter  of  water  weighs  one  gramme,  and  that 
a  liter  contains  1,000  cubic  centimeters.  A  liter  of  water 
therefore  weighs  1,000  grammes  or  one  kilogramme  (kilo.). 

For  conversion  into  other  systems  we  have: 

i  meter  =  39.37  inches. 

i  foot  =  0.304  meter. 

i  liter  =  61.03  cu-  m-  =  I-°6  U.  S.  qts. 

i  liter  =  33.81  U.  S.  fluidounces. 

i  gramme  =  15.43  grains. 

i  grain  =  0.0648  gramme. 

i  ounce  (apoth.)  =31.1  grammes. 

i  ounce  (avoirdupois)  =  28.35  grammes. 

i  pound  (apoth.)  =  373.2  grammes. 

i  pound  (avoirdupois)  =  453.6  grammes. 


INDEX 


Acetates,  112 
Acetic  acid,  112 

ether,  124 
Acetanilid,  135 
Acid,  17 

reaction,  17 
Acids,  free,  96 

reactions,  17,  96 

classification,  116 

determination,  163,  164 

identification,  114 

preliminary  tests,  114 

separation,  115 
Acidimetry,  158,  163 

standard  solutions  for,  160 
Acid  potassium,  tartrate  determina- 
tion, 165 

Alcohol,  ethyl  methyl,  124,  206 
Alkalimetry,  158,  164 

standard  solutions  for,  160 
Alkalies,  determination,  164,  166 

determination  of  compounds,  179 
Alkaline  reaction,  17 

earths,  27 

determination ,      177,     179, 
182,  183 

metals,  24 

hydrates   and   carbonates,    dis- 
tinction  209 
Alkaloids,  138 

general  reagents,  139 

separation,  138,  215 
Alloys,  analysis  of,  94 
Aluminum,  reactions,  36,  228 
Ammonium,  reactions,  25,  223 

hydroxid,  reagent,  206 

sulphid,  reagent,  206 
Amyl  alcohol,  126,  207 
Ammonia  in  water,  188,  199 
determination,  193 


Analysis  by  blow-pipe,  221 

by  neutralization,  158 

by  oxidation,  167 

by  precipitation,  177 

by  reduction,  167 
Animal  matter  in  water,  186,   192, 

iQ9 

Antidotes,  22 
Antifebrin,  135 
Antimony,  reactions,  71,  225 
Antipyrin,  135 
Arsenates,  reactions,  66 
Arsenic  acid,  reactions,  66 

reactions,  63,  225 

salts,  determination,  175 

in  milk,  219 

in  urine,  219 
Arsenites,  reactions,  65 
Arsenous  acid,  reactions,  65 

oxid,  64 
Arsine,  68 

Atomic  weights,  240 
Atropin,  143,  217 

Bacteria  in  water,  187,  199 
Barium,  reactions,  30,  223 
Base,  17 
Beakers,  3 
Benzene,  122,  207 
Benzine,  122 
Benzoic  acid,  130 
Benzol,  122 

Bettendorff's  test,  68,  73 
Bicarbonates,  99 
Bink's  burette,  150 
Bismuth,  reactions,  59,  225 
Blank  tests,  204 
Blast  lamp,  13 
Blow-pipe  flames,  15 
analysis,  221 


243 


244 


INDEX 


Blow-pipes,  15 

Boiler  scale,  186 

Boiling  point,  7 

Borax  bead,  226 

Boric  acid,  reactions,  101 

Bromids,  determination,  178,  180 

reactions,  108 

Bromin,  determination,  174 
Brucin,  142,  217 
Bunsen  burner,  12 

flame,  14 
Burettes,  150 

Cadmium,  reactions,  60,  226 
Calcium,  determination,  182 

reactions,  29,  223 

salts  in  water,  185,  187,  198 
Calculation  of  results,  148,  156 
Cane  sugar,  134 
Carbon,  223 

dioxid,  98 
Carbonates,  acid  and  normal,  98 

determination,  180 
Carbolic  acid,  127 
Carbonic  acid,  reactions,  98 
Cations,  separation  into  groups,  91 
Chemical  analysis,  kinds,  i 

incompatibilities,  21 

solutions,  i 

Citrates,  reactions,  113 
Citric  acid,  reactions,  113 
Chloral  hydrate,  129,  213 
Chlorates,  reactions,  in,  222 
Chloric  acids,  reactions,  in 
Chlorids,  determination,  178, 180,182 

reactions,  108 
Chlorin  in  water,  189,  195,  199 

determination,  174 
Chloroform,  122,  207,  213 
Chromates,  38,  100 
Chromic  acid,  reactions,  100 
Chromium,  reactions,  38,  227 

salts,  38 
Cobalt  glass,  24 


Cobalt  glass,  reactions,  44,  227 
Cocain,  144,  216 
Concentrated  solutions,  2 
Congo-red,  208 
Copper  in  coloring  matters,  219 

in  foods,  220 

in  water,  188 

reactions,  58,  227 
Creosote,  127 
Crucibles,  n 
Crystallization,  2 
Crystals,  2 
Cyanids,  reactions,  109 

determination,  178,  181 
Cylinders,  measuring,  150 

Decantation,  n 
Deoxidizing  flame,  16 
Desiccators,  7 
Dialysis,  210 
Dialyzers,  210 
Dilute  solutions,  2 
Distillation,  7,  220 
Dissociation,  18 

Elements,  240 

Empirical  standard  solutions,  154 

Erdmann's  float,  154 

Ethyl  acetate,  124 

alcohol,  124 
Evaporation,  5 
''Excess,"  1 6 

Ferric  compounds,  reactions,  42 
Ferricyanids,  reactions,  106 
Ferrocyanids,  reactions,  106 
Ferrous  compounds,  reactions,  41 

ammonium  sulphate,  standard 

solution,  169 
Filter-paper,  8,  207 

pump,  10 
Filtrate,  9 
Filtration,  8 
Flame,  Bunsen,  14 


INDEX 


245 


Flasks,  measuring,  150 
Fleitmann's  test  for  As,  68,  73 
Formaldehyde,  reactions,  128 
Fume  chamber,  7 
Furnace,  14 
Fusion,  ii 

Gallic  acid,  132 

Gastric  juice,  acidity,  164 

Glass,  solubility,  4 

Glucose,  reactions,  133 

Glycerin,  reactions,  1 26 

Gold,  reactions,  75,  228 

Grape  sugar,  133 

Graphite,  223 

Gravimetric  analysis,  148 

Griess'  test  for  organic  matter,  193, 

200 

Groups  of  acids,  116-118 
Group  I,  cations,  86,  89,  90 

II,  cations,  58,  77,  80 

III,  cations,  36,  45,  48 
Groups  III  and  IV,  cations,  separa- 
tion, 50,  55 

IV  and  V,  cations,  31,  33 
separation  of  metals  into,  91 
Gutzeit's  test  for  As,  and  Sb,  68,  73 

Hardness  of  water,  185,  199 

Heat,  sources  of,  12 

Hehner's  test  for  formaldehyd,  129 

Hood,  7,  212 

Hydriodic  acid,  reactions,  109 

Hydrobromic  acid,  reactions,  109 

Hydrochloric  acid,determination,i  79 

normal,  161 

reactions,  108 

as  reagent,  205 
Hydrocyanic  acid,  reactions,  109,  209 

determination,  181 

Hydroferricyanic  acid,  reactions,  106 

Hydroferrocyanic  acid,  reactions,  106 

Hydrogen  dioxid,  determination,  173 

sulphid,  reactions,  105 


Hydrogen  sulphid,  as  reagent,  206 

.     in  water,  191 
Hydrolysis,  20 
Hydrosulphuric  acid,  165 
Hydroxids,  determination,  179 
Hypochlorites,  reactions,  107 

determination,  174 
Hypochlorous  acid,  reactions,  107 
Hypophosphites,  reactions,  103 
Hypo  phosphoric  acid,  reactions,  103 
Hyposulphites,  reactions,  104 

Incompatibility,  21 
Indicators,  158,  159 
lodids,  determination,  178 

reactions,  109 
lodin,  determination,  174 

standard  solution,  170 
lodoform,  123 
Ionic  theory,  17 
Ions,  1 8 
Iron,  in  water,  189,  198 

determination,  169,  171 

reactions,  41,  227 

as  standard,  169 
Iron  salts,  41 

determination,  173,  175 
Isonitril,  123 
Isocyanid,  123 

Lactose,  134 

Lead,  reactions,  86,  226 

Lead  in  glazing,  220 

in  urine,  219 

in  water,  188,  220 
Liebig's  condenser,  212 
Litmus,  17 

Magnesium,  reactions,  27,  223 

in  water,  185,  187,  198 
Manganese,  reactions,  39,  226 
Marsh's  test  for  As,  68 
Mass  action,  18 
Measuring  apparatus,  150 


246 


INDEX 


Meconic  acid,  131,  218 
Melting  point,  7 
Meniscus,  153 

Mercuric  compounds,  determination, 
182 

reactions,  61 

Mercurous  reactions,  87 
Mercury,  reactions,  61,  225 

in  urine,  219 

Metaphosphates,  reactions,  103 
Metaphosphoric  acid,  reactions,  103 
Metallic  poisons,  218 
Metals,  separation  into  groups,  91 
Methyl  alcohol,  125 
Methyl  violet,  208 
Metric  system,  240 
Milk  sugar,  134 
Mineral  acids,  tests,  208 
Mohr's  burette,  150 
Morphin,  140,  217 
Mother-liquor,  3 

Narcotin,  141,  217 
Nickel,  reactions,  43,  227 
Nitrates,  reactions,  no 

determination,  179 

in  water,  190,  199 

determination,  196 
Nitric  acid,  reactions,  no,  208 
Nitrites,  determination,  195 

reactions,  107 

in  water,  189,  199 
Nitrous  acid,  reactions,  107 
Normal  solutions,  154 

Organic  compounds,  122 

matter  in  water,  186,  191,  197 

destruction,  218 
salts  of  alkalies,  determination, 

165,  166 

Crthophosphates,  reactions,  102 
Orthophosphoric  acid,  reactions,  102 
Oxalates,  reactions,  99 


Oxalic  acid,  99,  208,  217,  225 

normal,  162 
Oxidation,  15,  21 

analysis  by,  167 
Oxidimetry,  167 
Oxids,  determination,  180 
Oxidizing  flame,  15 

solutions,  standard,  167 


Petroleum  ether,  122,  207 
Pharmaceutical  incompatibility,  21 
Phenacetin,  136 
Phenol,  127,  213 

sulphonic  acid,  196 
Phosphates,  reactions,  102,  103 
Phosphorus,  detection,  213 

acids  of,  102,  103 
Physical  solution,  i 
Pipettes,  151 
Plaited  niters,  9 
Platinum  black,  76 

reactions,  76,  228 

substances  which  attack,  1 2 
Poisons,  203 
Porcelain  dishes,  4 
Potassium,  reactions,  24,  223 

chlorate,  as  reagent,  206 

cyanid,  determination,  181 

dichromate,  standard  solution, 
170 

permanganate,  standard  solu- 
tion, 167 

sulphocyanate  standard,  solu- 
tion, 178 
Precipitate,  8 
Precipitation,  8 

analysis  by,  177 
Prussian  blue,  43 
Purple  of  Cassius,  75 

Qualitative  analysis,  i 
Quantitative  analysis,  i 
Quinin,  145 


INDEX 


247 


Reactions,  i,  17 

Reagents,  i,  231 

Reducing  flame,  16 

Reduction,  15,  21 

Reinsch's  test  for  As  and  Sb,  67,  73 

Residual  titration,  179 

Reversible  reactions,  19 

Rubber  tubing,  207 

Saccharose,  134 
Salicylic  acid,  130 
Salol,  137 
Salts,  17 
Sand-bath,  3 
Saturated  solutions,  2 
Silicates,  223 
Silver,  reactions,  88,  228 

nitrate,  standard  solutions,  177, 

195 

Solids  in  water,  determination,  193 
Solute,  2 
Solutions,  i 

preparation  from  solids,  23 

theory  of,  17 
Sodium  reactions,  25,  223 

carbonate,  normal,  160 

chlorid,  determination,  180 

hydroxid,  normal,  160 
reagent,  206 

thiosulphate,  standard  solution, 

171 

Solubility  product,  20 
Sour  milk,  acidity,  164 
Standard  solutions,  154 

preparation,  158 

Stannic  compounds,  reactions,  74 
Stannous  compounds,  reactions,  74 
Starch,  133 
Steam-bath,  6 

Strontium,  reactions,  30,  223 
Strychnin,  141,  217 
Sugar,  cane,  134 

grape,  133 

milk,  134 


Sulphates,  reactions,  96,  222 

in  water,  189,  199 
Sulphids,  reactions,  105,  222 
Sulphites,  reactions,  97 
Sulphocyanates,  reactions,  106 
Sulphocyanic  acid,  106 
Sulphonal,  137 
Sulphur,  222 
Sulphuric  acid,  reactions,  96,  208 

as  reagent,  205 

Sulphurous  acid,  reactions,  97 
Supersaturated  solutions,  2 
Symbols  of  the  elements,  240 

Tannic  acid,  reactions,  131 

Tannins,  131 

Tartaric  acid,  reactions,  1 1 3 

Test-tubes,  3 

Tests  with  alloys  and  amalgams,  223, 

224 

Therapeutic  incompatibility,  21 
Thiosulphuric  acid,  104 
Tin,  reactions,  73,  228 
Toxicology.  203 
Trommer's  reaction^  123 

Vegetable  matter  in  water,  186,  191, 

200 

Veratrin,  144,  217 
Vinegar,  strength,  164 
Volumetric  analysis,  148 

Water,  detection,  222 

distilled,  204 

impurities,  185 

sanitary  examination,  185 

standards,  202 

bath,  6 
Washing,  10 

bottle,  10 

Zinc  in  water,  188 
reactions,  39,  228 


,    &. 


4Ifl 


